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Chemistry for Engineering Students

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Chemistry for Engineering Students

Lawrence S. Brown Texas A&M University

Thomas A. Holme Iowa State University

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SECOND EDITION

Chemistry for Engineering Students, Second Edition Lawrence S. Brown, Thomas A. Holme Publisher: Mary Finch Acquisitions Editor: Charles Hartford Developmental Editor: Rebecca Heider Assistant Editor: Ashley Summers

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Library of Congress Control Number: 2009935278 Student Edition: ISBN-13: 978-1-4390-4791-0 ISBN-10: 1-4390-4791-X Brooks/Cole 20 Davis Drive Belmont, CA 94002-3098 USA Cengage Learning is a leading provider of customized learning solutions with office locations around the globe, including Singapore, the United Kingdom, Australia, Mexico, Brazil, and Japan. Locate your local office at www.cengage.com/global.

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About the Authors Larry Brown is a Senior Lecturer and coordinator for the General Chemistry for Engineering Students course at Texas A&M University. He received his B.S. in 1981 from Rensselaer Polytechnic Institute, and his M.A. in 1983 and Ph.D. in 1986 from Princeton University. During his graduate studies, Larry spent a year working in what was then West Germany. He was a Postdoctoral Fellow at the University of Chicago from 1986 until 1988, at which time he began his faculty career at Texas A&M. Over the years, he has taught more than 10,000 general chemistry students, most of them engineering majors. Larry’s excellence in teaching has been recognized by awards from the Association of Former Students at Texas A&M at both the College of Science and University levels. A version of his class has been broadcast on KAMU-TV, College Station’s PBS affiliate. From 2001 to 2004, Larry served as a Program Officer for Education and Interdisciplinary Research in the Physics Division of the National Science Foundation. He also coordinates chemistry courses for Texas A&M’s engineering program in Doha, Qatar. When not teaching chemistry, he enjoys road bicycling and coaching his daughter Stephanie’s soccer team. Tom Holme is a Professor of Chemistry at Iowa State University and Director of the ACS Examinations Institute. He received his B.S. in 1983 from Loras College, and his Ph.D. in 1987 from Rice University. He began his teaching career as a Fulbright Scholar in Zambia, Africa and has also lived in Jerusalem, Israel and Suwon, South Korea. His research interests lie in computational chemistry, particularly as applied to understanding processes important for plant growth. He is also active chemical education research and has been involved with the general chemistry for engineers course at both Iowa State University and at the University of Wisconsin–Milwaukee where he was a member of the Chemistry and Biochemistry Department. He has received several grants from the National Science Foundation for work in assessment methods for chemistry, and the “Focus on Problem Solving” feature in this textbook grew out of one of these projects. He served as an Associate Editor on the encyclopedia “Chemistry Foundations and Applications.” In 1999 Tom won the ACS’s Helen Free Award for Public Outreach for his efforts doing chemical demonstrations on live television in the Milwaukee area.

v

Brief Contents 1

Introduction to Chemistry 1 OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

2

CLOSING INSIGHT THEME:


Stoichiometry


Gases


7

Air Pollution 126 Gas Sensors 148

Incandescent and Fluorescent Lights


Modern Light Sources: LEDs and Lasers

159 192

Chemical Bonding and Molecular Structure 200 Materials for Biomedical Engineering 201 Molecular Scale Engineering for Drug Delivery 234

Molecules and Materials 240 OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

Carbon 241 The Invention of New Materials

272

Energy and Chemistry 280 OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

vi

117

OPENING INSIGHT THEME:


9

Gasoline and Other Fuels 100 Alternative Fuels and Fuel Additives

The Periodic Table and Atomic Structure 158


8

91

125


6

Explosions 65 Explosives and Green Chemistry

99


5

Polymers 31 Polyethylene 56

Molecules, Moles, and Chemical Equations 64 OPENING INSIGHT THEME:

4

24

Atoms and Molecules 30 OPENING INSIGHT THEME:

3

Aluminum 2 Material Selection and Bicycle Frames

Energy Use and the World Economy Batteries 308

281

10

Entropy and the Second Law of Thermodynamics 318 CLOSING INSIGHT THEME:

Recycling of Plastics 319 The Economics of Recycling

Chemical Kinetics

347


Ozone Depletion 348 Tropospheric Ozone 379


11


12

Chemical Equilibrium OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

13

391

Concrete Production and Weathering Borates and Boric Acid 427

392

Electrochemistry 436 OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

14

335

Corrosion 437 Corrosion Prevention 465

Nuclear Chemistry 474 OPENING INSIGHT THEME: CLOSING INSIGHT THEME:

Cosmic Rays and Carbon Dating 475 Modern Medical Imaging Methods 498

Appendixes A B C D E F G H

International Table of Atomic Weights 507 Physical Constants 509 Electron Configurations of Atoms in the Ground State 510 Specific Heats and Heat Capacities of Some Common Substances Selected Thermodynamic Data at 298.15 K 512 Ionization Constants of Weak Acids at 25°C 518 Ionization Constants of Weak Bases at 25°C 520 Solubility Product Constants of Some Inorganic Compounds at 25°C 521 I Standard Reduction Potentials in Aqueous Solution at 25°C 523 J Answers to Check Your Understanding Exercises 526 K Answers to Odd-Numbered End-of-Chapter Exercises 529

511

Brief Contents

vii

Contents Preface

xix

Student Introduction

1

xxvii

Introduction to Chemistry 1 1.1 INSIGHT INTO Aluminum 1.2 The Study of Chemistry

2 4 6

1.3 The Science of Chemistry: Observations and Models Observations in Science 9 Interpreting Observations 10 Models in Science 11

1.4 Numbers and Measurements in Chemistry at Berkeley

National Laboratory, and the University of California

Courtesy of Zettl Research Group, Lawrence Berkeley

The Macroscopic Perspective 4 The Microscopic or Particulate Perspective Symbolic Representation 8

Units 13 Numbers and Significant Figures

16

1.5 Problem Solving in Chemistry and Engineering Using Ratios 18 Ratios in Chemistry Calculations 19 Conceptual Chemistry Problems 21 Visualization in Chemistry 22

1.6 INSIGHT INTO Material Selection and

Bicycle Frames 24 Focus on Problem Solving 25 Summary 26 Key Terms 26 Problems and Exercises 27

2

Atoms and Molecules 30 2.1 INSIGHT INTO Polymers

31 2.2 Atomic Structure and Mass 33 Fundamental Concepts of the Atom 33 Atomic Number and Mass Number 34 Isotopes 34 Atomic Symbols 35 Atomic Masses 36

2.3 Ions

38

Mathematical Description 38 Ions and Their Properties 39

viii

12 18

9

2.4 Compounds and Chemical Bonds

40

2.5 The Periodic Table

© Cengage Learning/Charles D. Winters

Chemical Formulas 40 Chemical Bonding 42

44

Periods and Groups 44 Metals, Nonmetals, and Metalloids

46

2.6 Inorganic and Organic Chemistry

47

Inorganic Chemistry—Main Groups and Transition Metals Organic Chemistry 49 Functional Groups 52

2.7 Chemical Nomenclature

48

53

Binary Systems 53 Naming Covalent Compounds 53 Naming Ionic Compounds 54

2.8 INSIGHT INTO Polyethylene

56

Focus on Problem Solving 58 Summary 59 Key Terms 59 Problems and Exercises 60

3

Molecules, Moles, and Chemical Equations 64 3.1 INSIGHT INTO Explosions

65 3.2 Chemical Formulas and Equations


67

Writing Chemical Equations 67 Balancing Chemical Equations 68

3.3 Aqueous Solutions and Net Ionic Equations Solutions, Solvents, and Solutes 72 Chemical Equations for Aqueous Reactions Acid–Base Reactions 78

72

76

3.4 Interpreting Equations and the Mole

81

Interpreting Chemical Equations 81 Avogadro’s Number and the Mole 82 Determining Molar Mass 83

3.5 Calculations Using Moles and Molar Masses

84

Elemental Analysis: Determining Empirical and Molecular Formulas Molarity 88 Dilution 90

3.6 INSIGHT INTO Explosives and Green Chemistry

86

91


4

Stoichiometry 99 4.1 INSIGHT INTO Gasoline and Other Fuels 4.2 Fundamentals of Stoichiometry

100

103

Obtaining Ratios from a Balanced Chemical Equation

104

Contents

ix

4.3 Limiting Reactants

108 4.4 Theoretical and Percentage Yields 113 4.5 Solution Stoichiometry 114 4.6 INSIGHT INTO Alternative Fuels and Fuel Additives 117 Focus on Problem Solving 118 Summary 119 Key Terms 119 Problems and Exercises 119

5

Gases 125 5.1 INSIGHT INTO Air Pollution Properties of Gases

5.2 Pressure

126

128

129

Measuring Pressure 130 Units of Pressure 130

5.3 History and Application of the Gas Law Units and the Ideal Gas Law

132

135

5.4 Partial Pressure

136 5.5 Stoichiometry of Reactions Involving Gases STP Conditions

139

140

5.6 Kinetic–Molecular Theory and Ideal Versus Real

Gases

141

Postulates of the Model 141 Real Gases and Limitations of the Kinetic Theory Correcting the Ideal Gas Equation 146

5.7 INSIGHT INTO Gas Sensors

145

148

Capacitance Manometer 148 Thermocouple Gauge 149 Ionization Gauge 149 Mass Spectrometer 151 Focus on Problem Solving 151 Summary 152 Key Terms 152 Problems and Exercises 152

6

The Periodic Table and Atomic Structure 158 6.1 INSIGHT INTO Incandescent and Fluorescent

Lights 159 6.2 The Electromagnetic Spectrum The Wave Nature of Light 161 The Particulate Nature of Light 165

6.3 Atomic Spectra The Bohr Atom

x

Contents

172

169

161

6.4 The Quantum Mechanical Model of the Atom Potential Energy and Orbitals Quantum Numbers 176 Visualizing Orbitals 178

173

175

6.5 The Pauli Exclusion Principle and Electron

Configurations

181

Orbital Energies and Electron Configurations Hund’s Rule and the Aufbau Principle 183

182

6.6 The Periodic Table and Electron

Configurations 185 6.7 Periodic Trends in Atomic Properties

187

Atomic Size 187 Ionization Energy 188 Electron Affinity 191

6.8 INSIGHT INTO Modern Light Sources: LEDs

and Lasers 192 Focus on Problem Solving 194 Summary 194 Key Terms 195 Problems and Exercises 195

7

Chemical Bonding and Molecular Structure 200 7.1 INSIGHT INTO Materials for Biomedical

Engineering 201 7.2 The Ionic Bond 202 Formation of Cations 202 Formation of Anions 204

7.3 The Covalent Bond

207

Chemical Bonds and Energy 207 Chemical Bonds and Reactions 209 Chemical Bonds and the Structure of Molecules

7.4 Electronegativity and Bond Polarity

209

211

Electronegativity 212 Bond Polarity 213

7.5 Keeping Track of Bonding: Lewis Structures Resonance

215

220

7.6 Orbital Overlap and Chemical Bonding

221

7.7 Hybrid Orbitals

224 7.8 Shapes of Molecules 226 7.9 INSIGHT INTO Molecular Scale Engineering for Drug Delivery 234 Focus on Problem Solving 235 Summary 235 Key Terms 236 Problems and Exercises 236

Contents

xi

8

Molecules and Materials 240 8.1 INSIGHT INTO Carbon

241

8.2 Condensed Phases—Solids

243 8.3 Bonding in Solids: Metals, Insulators, and Semiconductors 249 Models of Metallic Bonding 249 Band Theory and Conductivity 251 Semiconductors 252

8.4 Intermolecular Forces

256

Forces Between Molecules 256 Dispersion Forces 256 Dipole–Dipole Forces 258 Hydrogen Bonding 258

8.5 Condensed Phases—Liquids

261

Vapor Pressure 261 Boiling Point 263 Surface Tension 264

8.6 Polymers

265

Addition Polymers 266 Condensation Polymers 268 Copolymers 270 Physical Properties 271 Polymers and Additives 272

8.7 INSIGHT INTO The Invention

of New Materials

272


9

Energy and Chemistry 280 9.1 INSIGHT INTO Energy Use and the World

Economy 281 9.2 Defining Energy

284

Forms of Energy 284 Heat and Work 285 Energy Units 285

9.3 Energy Transformation and Conservation

of Energy

286

Waste Energy

288

9.4 Heat Capacity and Calorimetry Heat Capacity and Specific Heat Calorimetry 293

9.5 Enthalpy

295

Defining Enthalpy 295 DH of Phase Changes 296

xii

Contents

289

289

Vaporization and Electricity Production 298 Heat of Reaction 299 Bonds and Energy 299 Heats of Reaction for Some Specific Reactions 300

9.6 Hess’s Law and Heats of Reaction Hess’s Law 301 Formation Reactions and Hess’s Law

9.7 Energy and Stoichiometry Energy Density and Fuels

301

303

305

307

9.8 INSIGHT INTO Batteries

308


10

Entropy and the Second Law of Thermodynamics 318 10.1 INSIGHT INTO Recycling of Plastics 10.2 Spontaneity

319

320

Nature’s Arrow 320 Spontaneous Processes 321 Enthalpy and Spontaneity 321

10.3 Entropy

322

Probability and Spontaneous Change 322 Definition of Entropy 324 Judging Entropy Changes in Processes 324

10.4 The Second Law of Thermodynamics The Second Law 326 Implications and Applications

326

10.5 The Third Law of Thermodynamics 10.6 Gibbs Free Energy

326 327

330

Free Energy and Spontaneous Change Free Energy and Work 333

330

10.7 Free Energy and Chemical Reactions Implications of DG° for a Reaction

333

335

10.8 INSIGHT INTO The Economics of Recycling

335


11

Chemical Kinetics 347 11.1 INSIGHT INTO Ozone Depletion 11.2 Rates of Chemical Reactions Concept of Rate and Rates of Reaction

348 350 350

Contents

xiii

Stoichiometry and Rate 351 Average Rate and Instantaneous Rate

352

11.3 Rate Laws and the Concentration Dependence

of Rates

353

The Rate Law 354 Determination of the Rate Law

11.4 Integrated Rate Laws

355

358

Zero-Order Integrated Rate Law 359 First-Order Integrated Rate Law 360 Second-Order Integrated Rate Law 362 Half-Life 364

11.5 Temperature and Kinetics

366

Temperature Effects and Molecules That React Arrhenius Behavior 368

11.6 Reaction Mechanisms

366

373

Elementary Steps and Reaction Mechanisms 374 Mechanisms and Rate: The Rate-Determining Step

11.7 Catalysis

376

Homogeneous and Heterogeneous Catalysts Molecular Perspective of Catalysis 377 Catalysis and Process Engineering 379

Thomas A. Holme

376

11.8 INSIGHT INTO Tropospheric Ozone

377

379


12

Chemical Equilibrium 391 12.1 INSIGHT INTO Concrete Production and Weathering 12.2 Chemical Equilibrium

394

Forward and Reverse Reactions 394 Mathematical Relationships 397

12.3 Equilibrium Constants

398


The Equilibrium (Mass Action) Expression 398 Gas Phase Equilibria: Kp vs. Kc 399 Homogeneous and Heterogeneous Equilibria 400 Numerical Importance of the Equilibrium Expression 401 Mathematical Manipulation of Equilibrium Constants 402 Reversing the Chemical Equation 402 Adjusting the Stoichiometry of the Chemical Reaction 403 Equilibrium Constants for a Series of Reactions 404 Units and the Equilibrium Constant 405

xiv

Contents

12.4 Equilibrium Concentrations

405

Equilibrium Concentrations from Initial Concentrations 406 Mathematical Techniques for Equilibrium Calculations 409

12.5 LeChatelier’s Principle

410

Effect of a Change in Concentration of Reactant or Product on Equilibrium 410 Effect of a Change in Pressure on Equilibrium When Gases Are Present 412

392

Effect of a Change in Temperature on Equilibrium Effect of a Catalyst on Equilibrium 415

12.6 Solubility Equilibria

415

Solubility Product Constant 415 Defining the Solubility Product Constant 416 The Relationship Between Ksp and Molar Solubility Common Ion Effect 418 Reliability of Using Molar Concentrations 419

12.7 Acids and Bases

414

416

419

The Brønsted–Lowry Theory of Acids and Bases 420 The Role of Water in the Brønsted–Lowry Theory 420 Weak Acids and Bases 421

12.8 Free Energy and Chemical Equilibrium Graphical Perspective 425 Free Energy and Nonstandard Conditions

425

426

12.9 INSIGHT INTO Borates and Boric Acid

427


Electrochemistry 436 13.1 INSIGHT INTO Corrosion

437 13.2 Oxidation–Reduction Reactions and Galvanic Cells 438 Oxidation–Reduction and Half-Reactions 438 Building a Galvanic Cell 440 Terminology for Galvanic Cells 441 Atomic Perspective on Galvanic Cells 441 Galvanic Corrosion and Uniform Corrosion 442

13.3 Cell Potentials

444 Thomas A. Holme

13

Measuring Cell Potential 444 Standard Reduction Potentials 445 Nonstandard Conditions 449

13.4 Cell Potentials and Equilibrium Cell Potentials and Free Energy Equilibrium Constants 452

13.5 Batteries

453

Primary Cells 453 Secondary Cells 455 Fuel Cells 457 Limitations of Batteries

13.6 Electrolysis

450

450

457

458

Electrolysis and Polarity 458 Passive Electrolysis in Refining Aluminum 458 Active Electrolysis and Electroplating 460

13.7 Electrolysis and Stoichiometry

461

Current and Charge 461 Calculations Using Masses of Substances in Electrolysis 463 Contents

xv

13.8 INSIGHT INTO Corrosion Prevention Coatings 465 Cathodic Protection 466 Preventing Corrosion in Space

465

466


14

Nuclear Chemistry 474 14.1 INSIGHT INTO Cosmic Rays and Carbon Dating 14.2 Radioactivity and Nuclear Reactions

476

Radioactive Decay 476 Alpha Decay 477 Beta Decay 478 Gamma Decay 479 Electron Capture 479 Positron Emission 480

14.3 Kinetics of Radioactive Decay Radiocarbon Dating

481

483

14.4 Nuclear Stability

485 14.5 Energetics of Nuclear Reactions Binding Energy 487 Magic Numbers and Nuclear Shells

487

488

Lawrence S. Brown

14.6 Transmutation, Fission, and Fusion

489

Transmutation: Changing One Nucleus into Another Fission 490 Nuclear Reactors 492 Nuclear Waste 493 Fusion 494

14.7 The Interaction of Radiation and Matter Ionizing and Penetrating Power of Radiation Methods of Detecting Radiation 497 Measuring Radiation Dose 498

489

495

495

14.8 INSIGHT INTO Modern Medical Imaging

Methods

498


Appendixes A B C D

xvi

Contents

International Table of Atomic Weights 507 Physical Constants 509 Electron Configurations of Atoms in the Ground State Specific Heats and Heat Capacities of Some Common Substances 511

510

475

E F G H

Selected Thermodynamic Data at 298.15 K 512 Ionization Constants of Weak Acids at 25°C 518 Ionization Constants of Weak Bases at 25°C 520 Solubility Product Constants of Some Inorganic Compounds at 25°C 521 I Standard Reduction Potentials in Aqueous Solution at 25°C 523 J Answers to Check Your Understanding Exercises 526 K Answers to Odd-Numbered End-of-Chapter Exercises 529 Glossary Index

553

565

Contents

xvii


Preface The Genesis of This Text As chemists, we see connections between our subject and virtually everything. So the idea that engineering students should learn chemistry strikes most chemists as selfevident. But chemistry is only one of many sciences with which a practicing engineer must be familiar, and the undergraduate curriculum must find room for many topics. Hence, engineering curricula at more and more universities are shifting from the traditional year long general chemistry sequence to a single semester. And in most cases, these schools are offering a separate one-term course designed specifically for their engineering students. When schools—including our own—originally began offering these courses, there was no text on the market for them, so content from twosemester texts had to be heavily modified to fit the course. Although it is possible to do this, it is far from ideal. It became apparent that a book specifically geared for this shorter course was necessary. We have written this book to fill this need. Our goal is to instill an appreciation for the role of chemistry in many areas of engineering and technology and of the interplay between chemistry and engineering in a variety of modern technologies. For most engineering students, the chemistry course is primarily a prerequisite for courses involving materials properties. These courses usually take a phenomenological approach to materials rather than emphasizing the chemist’s molecular perspective. Thus one aim of this text is to provide knowledge of and appreciation for the chemical principles of structure and bonding that underpin materials science. This does not mean that we have written the book as a materials science text, but rather that the text is intended to prepare students for subsequent study in that area. The book also provides sufficient background in the science of chemistry for a technically educated professional. Engineering, after all, is the creative and practical application of a broad array of scientific principles, so its practitioners should have a broad base in the natural sciences.

Content and Organization The full scope of the traditional general chemistry course cannot be taught meaningfully in one semester or one or two quarters, so we have had to decide what content to include. There are basically two models used to condense the general chemistry curriculum. The first is to take the approach of an “essentials” book and reduce the depth of coverage and the number of examples but retain nearly all of the traditional topics. The second is to make more difficult and fundamental decisions as to what chemistry topics are proper and relevant to the audience, in this case future engineers. We chose the latter approach and built a 14-chapter book from the ground up to satisfy what we think are the goals of the course: • Provide a concise but thorough introduction to the science of chemistry. • Give students a firm foundation in the principles of structure and bonding as a foundation for further study of materials science. • Show the connection between molecular behavior and observable physical properties. • Show the connections between chemistry and the other subjects studied by engineering students, especially mathematics and physics.

xix

Taken together, the 14 chapters in this book probably represent somewhat more material than can comfortably fit into a standard semester course. Thus departments or individual instructors will need to make some further choices as to the content that is most suitable for their own students. We suspect that many instructors will not choose to include all of the material on equilibrium in Chapter 12, for example. Similarly, we have included more topics in Chapter 8, on condensed phases, than we expect most faculty will include in their courses.

Topic Coverage

Ames Laboratory

Courtesy of the U.S. Department of Energy’s

The coverage of topics in this text reflects the fact that chemists constantly use multiple concepts to understand their field, often using more than one model simultaneously. Thus the study of chemistry we present here can be viewed from multiple perspectives: macroscopic, microscopic, and symbolic. The latter two perspectives are emphasized in Chapters 2 and 3 on atoms, molecules, and reactions. In Chapters 4 and 5, we establish more of the connection between microscopic and macroscopic in our treatment of stoichiometry and gases. We return to the microscopic perspective to cover more details of atomic structure and chemical bonding in Chapters 6 through 8. The energetic aspects of chemistry, including important macroscopic consequences, are considered in Chapters 9 and 10, and kinetics and equilibrium are treated in Chapters 11 and 12, respectively. Chapter 13 deals with electrochemistry and corrosion, an important chemistry application for many engineering disciplines. Finally, we conclude with a discussion of nuclear chemistry.

Specific Content Coverage We know that there are specific topics in general chemistry that are vital to future engineers. We’ve chosen to treat them in the following ways. Organic Chemistry: Organic chemistry is important in many areas of engineering, particularly as related to the properties of polymers. Rather than using a single organic chapter, we integrate our organic chemistry coverage over the entire text, focusing on polymers. We introduce organic polymers in Section 2.1 and use polymers and their monomers in many examples in this chapter. Chapter 2 also contains a rich discussion of organic line structures and functional groups and ends with a section on the synthesis, structure, and properties of polyethylene. Chapter 4 opens and ends with discussions of fuels, a topic to which we return in Chapter 9. Chapter 8 contains more on carbon and polymers, and the recycling of polymers provides the context for consideration of the second law of thermodynamics in Chapter 10. Acid–Base Chemistry: Acid–base reactions represent another important area of chemistry with applications in engineering, and again we have integrated our coverage into appropriate areas of the text. Initially, we define acids and bases in conjunction with the introduction to solutions in Chapter 3. Simple solution stoichiometry is presented in Chapter 4. Finally, a more detailed treatment of acid-base chemistry is presented in the context of equilibria in Chapter 12. Nuclear Chemistry: A chapter dealing with nuclear chemistry, previously available as a custom option, has been added to the standard book for this edition. Coverage in this chapter includes fundamentals of nuclear reactions, nuclear stability and radioactivity, decay kinetics, and the energetic consequences of nuclear processes. Mathematics: The math skills of students entering engineering majors generally are stronger than those in the student body at large, and most of the students taking a course of the type for which this book is intended will be concurrently enrolled in

xx

Preface

an introductory calculus course. In light of this, we include references to the role of calculus where appropriate via our MathConnections boxes. These essays expand and review math concepts as they pertain to the particular topic being studied, and appear wherever the links between the topic at hand and mathematics seems especially strong. These boxes are intended to be unobtrusive, so those students taking a precalculus math course will not be adversely affected. The point of including calculus is not to raise the level of material being presented, but rather to show the natural connections between the various subjects students are studying.

Connections between Chemistry and Engineering

Lawrence S. Brown

Because this book is intended for courses designed for engineering majors, we strive to present chemistry in contexts that we feel will appeal to the interests of such students. Links between chemistry and engineering are central to the structure of the text. Each chapter begins and ends with a section called INSIGHT INTO . . . , which introduces a template or theme showing the interplay between chemistry and engineering. These sections are only the beginning of the connections, and the theme introduced in the initial Insight appears regularly throughout that chapter. We opt for currency in our engineering applications wherever possible, so throughout the book, we discuss recent key innovations in various fields. For example, Chapter 1 contains a brief discussion of OLEDs (organic light-emitting diodes), a new advance that appears likely to replace the liquid crystal screen in devices such as digital cameras and flat-panel computer monitors. OLEDs are revisited later in Chapter 6. In Chapter 2, we discuss the new polymer UHMWPE (ultra-high molecular weight polyethylene), which is stronger and lighter than Kevlar™ and is replacing Kevlar as filling in bulletproof vests. In Chapter 7, we describe mesoporous silicon nanoparticles, a front-line research topic that may have important applications in biomedical engineering in the future.

Approach to Problem Solving Problem solving is a key part of college chemistry courses and is especially important as a broadly transferable skill for engineering students. Accordingly, this text includes worked problems throughout. All of our Example Problems include a Strategy section immediately following the problem statement, in which we emphasize the concepts and relationships that must be considered to work the problem. After the solution, we often include a section called Analyze Your Answer that is designed to help students learn to estimate whether or not the answer they have obtained is reasonable. In many examples, we also include Discussion sections that help explain the importance of a problem solving concept or point out common pitfalls to be avoided. Finally, each example closes with a Check Your Understanding problem or question to help the student to generalize or extend what’s been learned in the example problem. We believe that the general chemistry experience should help engineering students develop improved problem solving skills. Moreover, we feel that those skills should be transferable to other subjects in the engineering curriculum even though chemistry content may not be involved. Accordingly, we include a unique feature at the end of each chapter called FOCUS ON PROBLEM SOLVING. In these sections, the questions posed do not require a numerical answer, but rather ask the student to identify the strategy or reasoning to be used in the problem and often require them to

Preface

xxi

identify missing information for the problem. In most cases, it is not possible to arrive at a final numerical answer using the information provided, so students are forced to focus on developing a solution rather than just identifying and executing an algorithm. The end-of-chapter exercises include additional problems of this nature so the Focus on Problem Solving can be fully incorporated into the course. This feature grew out of an NSF-funded project on assessing problem solving in chemistry classes.

Text Features We employ a number of features, some of which we referred to earlier, to help students see the utility of chemistry and understand the connections to engineering.

INSIGHT INTO Sections Each chapter is built around a template called Insights Into . . . . These themes, which both open and close each chapter, have been chosen to showcase connections between engineering and chemistry. In addition to the chapter opening and closing sections, the template themes are woven throughout the chapter, frequently providing the context for points of discussion or example problems. This special Insight icon is used throughout the book to identify places where ideas presented in the chapter opening section are revisited in the narrative.

FOCUS ON PROBLEM SOLVING Sections Engineering faculties unanimously say that freshman engineering students need practice in solving problems. However, it is important to make a distinction here between problems and exercises. Exercises provide a chance to practice a narrow skill, whereas problems require multiple steps and thinking outside the context of the information given. Focus on Problem Solving offers students the chance to develop and practice true problem solving skills. These sections, which appear at the end of every chapter, include a mix of quantitative and qualitative questions that focus on the process of finding a solution to a problem, not the solution itself. We support these by including additional similar problems in the end-of-chapter material. MathConnections In our experience, one trait that distinguishes engineering students from other general chemistry students is a higher level of comfort with mathematics. Typically most students who take a class of the sort for which this book has been written will also be taking a course in calculus. Thus it seems natural to us to point out the mathematical underpinnings of several of the chemistry concepts presented in the text because this should help students forge mental connections between their courses. At the same time, we recognize that a student taking a precalculus math course should not be precluded from taking chemistry. To balance these concerns, we have placed any advanced mathematics into special MathConnections sections, which are set off from the body of the text. Our hope is that those students familiar with the mathematics involved will benefit from seeing the origin of things such as integrated rate laws, whereas those students with a less extensive background in math will still be able to read the text and master the chemistry presented. Example Problems Our examples are designed to illustrate good problem solving practices by first focusing on the reasoning behind the solution before moving into any needed calculations. We emphasize this “think first” approach by beginning with a Strategy section, which outlines a plan of attack for the problem. We find that many students are too quick to accept whatever answer their calculator might display. To combat this, we follow most solutions with an Analyze Your Answer section, which uses estimation and other strategies to walk students through a double check of their answers. Every example closes with a Check Your Understanding exercise to allow students to practice or extend the skill they have just learned. Answers to these additional exercises are included in Appendix J at the end of the book. End-of-Chapter Features Each chapter concludes with a chapter summary, outlining the main points of the chapter, and a list of key terms, each of which includes the

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section number where the term first appeared. Definitions for all key terms appear in the Glossary. Problem Sets Each chapter includes roughly 100 problems and exercises, spanning a wide range of difficulty. Most of these exercises are identified with specific sections to provide the practice that students need to master material from that section. Each chapter also includes a number of Additional Problems, which are not tied to any particular section and which may incorporate ideas from multiple sections. Focus on Problem Solving exercises follow, as described earlier. The problems for most chapters conclude with Cumulative Problems, which ask students to synthesize information from the current chapter with what they’ve learned from previous chapters to form answers. Answers for all odd-numbered problems appear at the end of the book in Appendix K. Margin Notes Margin notes in the text point out additional facts, further emphasize points, or point to related discussion either earlier or later in the book.

New in this Edition There are several key changes in this second edition of the textbook. In addition to being able to catch and fix minor errors from the first edition, we were also able to find out which of the “Insight Into. . .” sections were the least successful at engaging student interest for that subject. Thus, we have introduced two new topics for the chapter-opening insights: Materials for Biomedical Engineering in Chapter 7 and Concrete Production and Weathering in Chapter 12. Both of these themes are more readily connected to engineering applications than those that they replaced from the first edition. The closing insight sections for Chapters 3 & 7 have also been rewritten to highlight topics with more current relevance. Because we realize that some instructors wish to include the topic in their courses, this edition includes a final chapter dealing with nuclear chemistry. We have also made significant changes to the end of chapter problems throughout the book. Approximately 25% of the problems in this edition are new, with most of the changes focused on two objectives. First, the new edition is integrated with the OWL electronic homework system, and a sizable majority of the new problems are available in OWL. This will make it significantly easier for instructors who would like to use OWL in their classes to achieve a strong correlation between problem assignments and the text. Second, we have worked to add a number of new problems that have a strong engineering focus. This addition is designed to provide more emphasis on the connections between the chemistry topics in this book and the engineering careers that the students who read it are pursuing. Many of these engineering problems are also available in OWL.

Supplements for the Instructor Faculty Companion Website Accessible from www.cengage.com/chemistry/brown, this website provides WebCT and Blackboard versions of ExamView Computerized Testing.

Instructor’s Resource CD-DVD Package ISBN-10: 1-439-04982-3 This collection of book-specific lecture and class tools is the fastest and easiest way to build powerful, customized, media-rich lectures. The CD includes chapter-specific PowerPoint Lecture presentations, a library of images from the text, the Instructor Solutions Manual, and sample chapters from the Student Solutions Manual and Study Guide. Also included are JoinIn™ questions for Response Systems, which let

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you transform your classroom and assess your students’ progress with instant in-class quizzes and polls. The Chemistry Multimedia Library DVD contains lecture-ready animations, simulations, and movies.

ExamView® Computerized Testing CD-ROM ISBN-10: 0-538-73523-6 Featuring automatic grading, EXAMVIEW allows you to create, deliver, and customize tests and study guides (both print and online) in minutes using the questions from the book’s test bank. See assessments onscreen exactly as they will print or display online. Build tests of up to 250 questions using up to 12 question types and enter an unlimited number of new questions or edit existing questions.

Supplements for the Student Student Solutions Manual and Study Guide by Steve Rathbone of Blinn College ISBN-10: 1-439-04981-5 The STUDENT SOLUTIONS MANUAL AND STUDY GUIDE provides students with a comprehensive guide to working the solutions to the odd-numbered end-of-chapter problems in the text and also includes each chapter’s Study Goals and Chapter Objective quizzes. Because the best way for students to learn and understand the concepts is to work multiple, relevant problems on a daily basis and to have reinforcement of important topics and concepts from the book, the STUDENT SOLUTIONS MANUAL gives students instant feedback by providing not only the answers to problems, but also detailed explanations of each problem’s solution.

OWL for General Chemistry OWL Instant Access (1 Semester) ISBN-10: 0-495-05098-9 e-Book in OWL Instant Access (1 Semester) ISBN-10: 0-538-73313-6 Authored by Roberta Day and Beatrice Botch of the University of Massachusetts, Amherst, and William Vining of the State University of New York at Oneonta. OWL includes more assignable, gradable content (including end-of-chapter questions specific to this textbook), more reliability, and more flexibility than any other system. Developed by chemistry instructors for teaching chemistry, OWL makes homework management a breeze and has already helped hundreds of thousands of students master chemistry through tutorials, interactive simulations, and algorithmically generated homework questions that provide instant, answer-specific feedback. In addition, OWL users (instructors and students) experience service that goes far beyond the ordinary. OWL is continually enhanced with online learning tools to address the various learning styles of today’s students such as: • e-Books, which offer a fully integrated electronic textbook correlated to OWL questions • Go Chemistry® mini video lectures • Quick Prep review courses that help students learn essential skills to succeed in General and Organic Chemistry • Thinkwell Video Lessons that teach key concepts through video, audio, and whiteboard examples • Jmol molecular visualization program for rotating molecules and measuring bond distances and angles • Parameterized end-of-chapter questions designed specifically to match this text xxiv

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Go Chemistry® for General Chemistry (27-module set) ISBN-10: 1-439-04700-6 GO CHEMISTRY® is a set of 27 exceptional mini video lectures on essential chemistry topics that students can download to their video iPod, iPhone, or portable video player—ideal for the student on the go! Developed by award-winning chemists, these new electronic tools are designed to help students quickly review essential chemistry topics. Mini video lectures include animations and problems for a quick summary of key concepts. Selected modules include e-flashcards that briefly introduce key concepts and then test student understanding of the basics with a series of questions. GO CHEMISTRY also plays on QuickTime, iTunes, and Windows Media Player. For a complimentary look at the modules, visit www.cengage.com/go/chemistry where you can view and download two demo modules.

Acknowledgments We are very excited to see this book move forward in this second edition, and we are grateful for the help and support we have enjoyed from a large and talented team of professionals. There are many people without whom we never could have done this. Foremost among them are our families, to whom this book is again dedicated. The origin of this text can be traced back many years, and a long list of people at Brooks/Cole played important roles. Jennifer Laugier first brought the two of us together to work on a book for engineering students. Jay Campbell’s work as developmental editor for the first edition was tremendous, and without his efforts the book might never have been published. When Jay became involved, the project had been languishing for some time, and the subsequent gains in momentum were clearly not coincidental. The editorial leadership team at that time, consisting of Michelle Julet, David Harris, and Lisa Lockwood, was also crucial in seeing this project come to fruition. The decision to launch a book in a market segment that has not really existed was clearly not an easy one, and we appreciate the confidence that everyone at Brooks/Cole placed in us. Like any modern business, the publishing industry seems to be one of constant change. Perhaps most obviously, our publisher is now known as Brooks/Cole Cengage Learning. And as we set out to work on this second edition, a number of changes had taken place in the Brooks/Cole team. Charlie Hartford and Lisa Lockwood supported us in the decision to go forward with this new edition and contributed valuable ideas leading to what we believe are substantial improvements. As our new developmental editor, Rebecca Heider has seen us through the entire revision process. When things were running behind schedule, she helped get us back on track. Lisa Weber has coordinated the integration of our text with the OWL homework system, this integration being one of the major undertakings of this revision. Teresa Trego managed the actual production process, and most of the production work was done by Pre-Press PMG under the leadership of Patrick Franzen. Within Pre-Press, a talented team of individuals has handled all aspects of production, including copyediting, illustration, photo research, and page layout. Allen Apblett has again served as an accuracy checker during the page proof stage of production. Jon Olafsson has overseen revisions to ancillary materials. The book in your hands truly reflects the best efforts of many hard working professionals, and we are grateful to all of them for their roles in this project. In preparing the new material for this edition, we have also been helped by colleagues with expertise in specific areas. Conversations with Victor Lin and Klaus Schmidt-Rohr of Iowa State University led to the development of the new insight sections involving biomaterials and concrete. Sherry Yennello of Texas A&M University provided much needed advice and assistance with the nuclear chemistry chapter. It has been nearly four years since the first edition was published, and over that time we have received useful feedback from numerous students and colleagues. Much of that feedback was informal, including e-mail from students or faculty members Preface

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pointing out errors they have found or letting us know about sections they really liked. Although there is no way to list all of the people who have contributed in this way, we do sincerely thank you all. Faculty members from a wide variety of institutions also provided more formal comments on the text at various stages of its development. We thank the following reviewers for their contributions to the current revision. Paul A. DiMilla, Northeastern University Walter England, University of Wisconsin–Milwaukee Mary Hadley, Minnesota State University, Mankato Andy Jorgensen, University of Toledo Karen Knaus, University of Colorado–Denver Pamela Wolff, Carleton University Grigoriy Yablonsky, Saint Louis University We also thank the following reviewers for their contributions to the development of the first edition of the book. Robert Angelici, Iowa State University Allen Apblett, Oklahoma State University Jeffrey R. Appling, Clemson University Rosemary Bartoszek-Loza, The Ohio State University Danny Bedgood, Charles Sturt University James D. Carr, University of Nebraska Victoria Castells, University of Miami Paul Charlesworth, Michigan Technological University Richard Chung, San Jose State University Charles Cornett, University of Wisconsin—Platteville Robert Cozzens, George Mason University Ronald Evilia, University of New Orleans John Falconer, University of Colorado Sandra Greer, University of Maryland Benjamin S. Hsaio, State University of New York at Stony Brook Gerald Korenowski, Rensselaer Polytechnic Institute Yinfa Ma, University of Missouri—Rolla Gerald Ray Miller, University of Maryland Linda Mona, Montgomery College Michael Mueller, Rose-Hulman Institute of Technology Kristen Murphy, University of Wisconsin—Milwaukee Thomas J. Murphy, University of Maryland Richard Nafshun, Oregon State University Scott Oliver, State University of New York at Binghamton The late Robert Paine, Rochester Institute of Technology Steve Rathbone, Blinn College Jesse Reinstein, University of Wisconsin—Platteville Don Seo, Arizona State University Mike Shaw, Southern Illinois University—Edwardsville Joyce Solochek, Milwaukee School of Engineering Jack Tossell, University of Maryland Peter T Wolczanski, Cornell University Larry Brown Tom Holme October, 2009

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Student Introduction Chemistry and Engineering As you begin this chemistry course, odds are that you may be wondering “Why do I have to take chemistry anyway? I’ll never really need to know any of this to be an engineer.” So we’d like to begin by offering just a few examples of the many links between our chosen field of chemistry and the various branches of engineering. The most obvious examples, of course, might come from chemical engineering. Many chemical engineers are involved with the design or optimization of processes in the chemical industry, so it is clear that they would be dealing with concepts from chemistry on a daily basis. Similarly, civil or environmental engineers working on environmental protection or remediation might spend a lot of time thinking about chemical reactions taking place in the water supply or the air. But what about other engineering fields? Much of modern electrical engineering relies on solid-state devices whose properties can be tailored by carefully controlling their chemical compositions. And although most electrical engineers do not regularly make their own chips, an understanding of how those chips operate on an atomic scale is certainly helpful. As the push for ever smaller circuit components continues, the ties between chemistry and electrical engineering will grow tighter. From organic light-emitting diodes (OLEDs) to single molecule transistors, new developments will continue to move out of the chemistry lab and into working devices at an impressive pace. Some applications of chemistry in engineering are much less obvious. At 1483 feet, the Petronas Towers in Kuala Lumpur, Malaysia, were the tallest buildings in the world when they were completed in 1998. Steel was in short supply in Malaysia, so the towers’ architects decided to build the structures out of something the country had an abundance of and local engineers were familiar with: concrete. But the impressive height of the towers required exceptionally strong concrete. The engineers eventually settled on a material that has come to be known as high strength concrete, in which chemical reactions between silica fume and portland cement produce a stronger material, more resistant to compression. This example illustrates the relevance of chemistry even to very traditional fields of engineering, and we will discuss some aspects of the chemistry of concrete in Chapter 12.

About This Text Both of us have taught general chemistry for many years, and we are familiar with the difficulties that students may encounter with the subject. Perhaps more importantly, for the past several years, we’ve each been teaching engineering students in the type of one semester course for which this text is designed. The approach to subjects presented in this text draws from both levels of experience. We’ve worked hard to make this text as readable and student friendly as possible. One feature that makes this book different from any other text you could have used for this course is that we incorporate connections between chemistry and engineering as a fundamental component of each chapter. You will notice that each chapter begins and ends with a section called INSIGHT INTO. . . . These sections are only the beginning of the connections, and the theme introduced in the initial insight appears regularly throughout that chapter. This special icon identifies material that is closely related to the theme of the chapter opening Insight section. We’ve heard many

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students complain that they don’t see what chemistry has to do with their chosen fields, and we hope that this approach might help you to see some of the connections. Engineering students tend to take a fairly standard set of courses during their first year of college, so it’s likely that you might be taking calculus and physics courses along with chemistry. We’ve tried to point out places where strong connections between these subjects exist, and at the same time to do this in a way that does not disadvantage a student who might be taking a precalculus math class. Thus we may refer to similarities between equations you see here and those you might find in a physics text, but we do not presume that you are already familiar with those equations. In the case of math, we use special sections called MathConnections to discuss the use of math, and especially calculus, in chemistry. If you are familiar with calculus or are taking it concurrently with this class, these sections will help you to see how some of the equations used in chemistry emerge from calculus. But if you are not yet taking calculus, you can simply skip over these sections and still be able to work with the needed equations. Although our primary intent is to help you learn chemistry, we also believe that this text and the course for which you are using it can help you to develop a broad set of skills that you will use throughout your studies and your career. Foremost among them is problem solving. Much of the work done by practicing engineers can be characterized as solving problems. The problems you will confront in your chemistry class clearly will be different from those you will see in engineering, physics, or math. But taken together, all of these subjects will help you formulate a consistent approach that can be used to attack virtually any problem. Many of our students tend to “jump right in” and start writing equations when facing a problem. But it is usually a better idea to think about a plan of attack before doing that, especially if the problem is difficult or unfamiliar. Thus all of our worked examples include a Strategy section in which we outline the path to a solution before starting to calculate anything. The Solution section then puts that strategy into action. For most numerical examples, we follow the solution with a section we call Analyze Your Answer, in which we use estimation or comparison to known values to confirm that our result makes sense. We’ve seen many students who believe that whatever their calculator shows must be the right answer, even when it should be easy to see that a mistake has been made. Many examples also include a Discussion section in which we might talk about common pitfalls that you should avoid or how the problem we’ve just done relates to other ideas we’ve already explored. Finally, each example problem closes with a Check Your Understanding question or problem, which gives you a chance to practice the skills illustrated in the example or to extend them slightly. Answers to these Check Your Understanding questions appear in Appendix J. While we are thinking about the example problems, a few words about rounding and significant figures are in order. In solving the example problems, we have used atomic weights with the full number of significant figures shown in the Periodic Table inside the back cover. We have also used as many significant figures as available for constants such as the speed of light or the universal gas constant. Where intermediate results are shown in the text, we have tried to write them with the appropriate number of significant figures. But when those same intermediate results are used in a subsequent calculation, we have not rounded the values. Instead we retain the full calculator result. Only the final answer has actually been rounded. If you follow this same procedure, you should be able to duplicate our answers. (The same process has been used to generate the answers to numerical problems appearing in Appendix K.) For problems that involve finding the slope or intercept of a line, the values shown have been obtained by linear regression using the algorithms built into either a spreadsheet or a graphing calculator. A unique feature of this text is the inclusion of a Focus on Problem Solving question at the end of each chapter. These questions are designed to force you to think about the process of solving the problem rather than just getting an answer. In many cases, these problems do not include sufficient information to allow you to reach a

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final solution. Although we know from experience that many beginning engineering students might find this frustrating, we feel it is a good approximation to the kind of problems that a working engineer might confront. Seldom would a client sit down and provide every piece of information that you need to solve the problem at hand. One of the most common questions we hear from students is “How should I study for chemistry?” Sadly, that question is most often asked after the student has done poorly on one or more exams. Because different people learn best in different ways, there isn’t a single magic formula to ensure that everyone does well in chemistry. But there are some common strategies and practices that we can recommend. First and foremost, we suggest that you avoid getting behind in any of your classes. Learning takes time, and very few people can master three chapters of chemistry (or physics, or math, or engineering) the night before a big exam. Getting behind in one class inevitably leads to letting things slide in others, so you should strive to keep up from the outset. Most professors urge students to read the relevant textbook material before it is presented in class. We agree that this is the best approach, because even a general familiarity with the ideas being presented will help you to get a lot more out of your class time. In studying for exams, you should try to make a realistic assessment of what you do and don’t understand. Although it can be discomforting to focus on the problems that you don’t seem to be able to get right, spending more time studying things that you have already mastered will probably have less impact on your grade. Engineering students tend to focus much of their attention on numerical problems. Although such calculations are likely to be very important in your chemistry class, we also encourage you to try to master the chemical concepts behind them. Odds are that your professor will test you on qualitative or conceptual material, too. Finally, we note that this textbook is information rich. It includes many of the topics that normally appear in a full year college chemistry course, but it is designed for a course that takes only one semester. To manage the task of paring down the volume of materials, we’ve left out some topics and shortened the discussion of others. Having the Internet available means that you can always find more information if what you have read sparks your interest. We are excited that this book has made it into your hands. We hope you enjoy your semester of learning chemistry and that this book is a positive part of your experience. Larry Brown Tom Holme October, 2009


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2 1

Introduction Chapter Titleto Chemistry

OUTLINE 1.1

INSIGHT INTO Aluminum

1.2

The Study of Chemistry

1.3

The Science of Chemistry: Observations and Models

1.4

Numbers and Measurements in Chemistry

1.5

Problem Solving in Chemistry and Engineering

1.6

INSIGHT INTO Material Selection and Bicycle Frames

Scientists from Lawrence Berkeley National Laboratory and the University of California at Berkeley developed this nanoscale “conveyor belt.” Individual metal atoms are transported along a carbon nanotube from one metal droplet to another. This research offers a possible means for the atomic scale construction of optical, electronic, and mechanical devices. Courtesy of Zettl Research Group, Lawrence Berkeley National Laboratory, and the University of California at Berkeley

I

n the not too distant future, engineers may design and assemble miniature mechanical or electronic devices, gears, and other parts fabricated on an atomic scale. Their decisions will be guided by knowledge of the sizes and properties of the atoms of different elements. Such devices might be built up atom by atom: each atom would be specified based on relevant design criteria and maneuvered into position using techniques such as the “conveyor belt” shown above. These nanomachines will be held together not by screws or rivets but by the forces of attraction between the different atoms—by chemical bonds. Clearly, these futuristic engineers will need to understand atoms and the forces that bind them together. In other words, they will need to understand chemistry. At least for now, though, this atomic level engineering remains in the future. But what about today’s practicing engineers? How do their decisions depend on knowledge of chemistry? And from your own perspective as an engineering student, why are you required to take chemistry?

Nanoscience deals with objects whose sizes are similar to those of atoms and molecules. Try a web search for “nanoscience” or “molecular machines” to learn more.

Online homework for this Online for this chapterhomework may be assigned chapter may be assigned in OWL in OWL.

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The Accreditation Board for Engineering and Technology, or ABET, is a professional organization that oversees engineering education. According to ABET’s definition, “Engineering is the profession in which a knowledge of the mathematical and natural sciences gained by study, experience, and practice is applied with judgment to develop ways to utilize, economically, the materials and forces of nature for the benefit of mankind.” So as one of the sciences, chemistry is clearly included in the realm of knowledge at the disposal of an engineer. Yet engineering students do not always recognize the role of chemistry in their chosen profession. One of the main goals of this textbook is to instill an appreciation of the role of chemistry in many areas of engineering and technology and in the interplay between chemistry and engineering in a variety of modern technologies. The study of chemistry involves a vast number of concepts and skills. The philosophy of this book is to present those basic ideas and also to apply them to aspects of engineering where chemistry is important. Each chapter will begin with an example of chemistry related to engineering. Some of these examples, such as the burning of fuels, will involve fairly clear applications of chemical principles and reactions. In other cases, the role of chemistry may be less immediately apparent. In Chapter 6, we will consider how evolving knowledge of chemical properties has driven the design of different light sources, from the simple incandescent bulb through modern lasers and organic light emitting diodes (OLEDs). Other themes will involve the design and selection of materials for various uses and the importance of chemistry in environmental engineering problems. All of these chapter-opening sections have titles that begin with “Insight into . . . ,” and the questions that are raised in them will guide our exploration of the relevant fundamentals of chemistry presented throughout that chapter. Our first case considers the production of aluminum and the history of aluminum as a structural material.

Chapter Objectives After mastering this chapter, you should be able to ❚ describe how chemistry and engineering helped transform aluminum from a precious metal into an inexpensive structural material. ❚ explain the usefulness of the macroscopic, microscopic, and symbolic perspectives in understanding chemical systems. ❚ draw pictures to illustrate simple chemical phenomena (like the differences among solids, liquids, and gases) on the molecular scale. ❚ explain the difference between inductive and deductive reasoning in your own words. ❚ use appropriate ratios to convert measurements from one unit to another. ❚ express the results of calculations using the correct number of significant figures.

INSIGHT INTO

1.1 Each year, more than 100 billion aluminum cans are produced in the United States.

2

Aluminum

If you are thirsty, you might ask yourself several questions about what to drink. But you probably wouldn’t ask, “Where did the can that holds this soda come from, and why is it made of aluminum?” The aluminum can has become so common that it’s easy to take for granted. What makes aluminum an attractive material for this type of application, and how did it become such a familiar part of life?

Chapter 1 Introduction to Chemistry

You probably can identify a few properties of aluminum that make it suitable for use in a soda can. Compared with most other metals, aluminum is light but fairly strong. So a typical aluminum can is much lighter than a comparable tin or steel can. This means that the can does not add much weight compared to the soda itself, so the cans are easier to handle and cheaper to ship. A soda can made of lead certainly would be less convenient. The fact that aluminum does not readily undergo chemical reactions that might degrade it as the cans are transported and stored is also important. But although all of those features of the aluminum can are nice, they wouldn’t be of much practical use if aluminum were not readily available and reasonably inexpensive. The widespread availability of aluminum results from an impressive collaboration between the basic science of chemistry and the applied sciences of engineering. In the 19th century, aluminum was a rare and precious material. In Europe, Napoleon was emperor of a sizable portion of the continent, and he would impress guests by using extravagant aluminum tableware. In the United States, when architects wanted a suitably impressive material for the capstone at the top of the Washington Monument, a tribute to the “father of our country,” they chose aluminum. Weighing in at 100 ounces, the capstone of the monument was the largest single piece of pure aluminum ever cast at that time. Yet today, sheets of aluminum weighing more than 100 pounds are regularly found in many metal shops. Why was aluminum so expensive then, and what changed to make it so affordable now? Initial discussion of this question can be framed in terms of Figure 1.1, which looks rather broadly at the interactions of human society with the earth. Society, represented by the globe, has needs for goods and materials. Currently, and for the foreseeable future, the raw materials needed to make these goods must somehow be extracted from the earth. When the goods are used up, the leftovers become waste that must be disposed of, completing the cycle by returning the exhausted materials to the ecosystem. Ultimately, the role of engineering in this cycle is to maximize the efficiency with which materials are extracted and minimize the amount of waste that is returned. Figure 1.1 ❚ The interactions of human society with the earth can be thought of largely in terms of the conversion of matter from raw materials into waste. Much of engineering consists of efforts to optimize the processes used in these conversions. And as the science of matter, chemistry is an important element of the knowledge exploited in engineering those processes.

Matter flows from the ecosphere into the human economy as raw materials.

Human society

Ecosphere

Matter flows from the human economy into the ecosphere as waste.

1.1

Aluminum

3

The aluminum in bauxite is typically found in one of three minerals: gibbsite, bohmite, and diaspore.

Let’s think about aluminum in this context. Pure aluminum is never found in nature. Instead, the metal occurs in an ore, called bauxite, that is composed of both useless rock and aluminum in combination with oxygen. So before aluminum can be used in our soda can, it must first be extracted or “won” from its ore and purified. Because aluminum combines very readily with oxygen, this presents some serious challenges. Some of these challenges are chemical and will be revisited in Chapter 13 of this text. Some of the early steps, however, can be solved by clever applications of physical properties, and we will consider a few of them as we investigate introductory material in this chapter. When confronted with a complex mixture of materials, such as an ore, how does a chemist think about separating the mixture? To look into this type of question, we should adopt the approach that is commonly taken in science. The term scientific method has various possible definitions, and we’ll look at this concept further in Section 1.4. But at this point, we will consider it an approach to understanding that begins with the observation of nature, continues to hypothesis or model building in response to that observation, and ultimately includes further experiments that either bolster or refute the hypothesis. In this definition, the hypothesis is an educated guess at ways to explain nature. In this chapter, we will see how this method relates to chemistry in general and also to issues relating to materials like aluminum and their use in society.

1.2

The Study of Chemistry

Chemistry has been called the “central science” because it is important to so many other fields of scientific study. So, even if you have never taken a chemistry course, chances are good that you have seen some chemistry before. This text and the course in which you are using it are designed to help you connect pieces of information you have already picked up, increase your understanding of chemical concepts, and give you a more coherent and systematic picture of chemistry. The ultimate goal of introductory college chemistry courses is to help you appreciate the chemical viewpoint and the way it can help you to understand the natural world. This type of perspective of the world is what enables chemists and engineers to devise strategies for refining metals from their ores, as well as to approach the many other applied problems we’ll explore. This coherent picture involves three levels of understanding or perspectives on the nature of chemistry: macroscopic, microscopic, and symbolic. By the end of this course, you should be able to switch among these perspectives to look at problems involving chemistry in several ways. The things we can see about substances and their reactions provide the macroscopic perspective. We need to interpret these events considering the microscopic (or “particulate”) perspective, where we focus on the smallest components of the system. Finally, we need to be able to communicate these concepts efficiently, so chemists have devised a symbolic perspective that allows us to do that. We can look at these three aspects of chemistry first, to provide a reference for framing our studies at the outset.

The Macroscopic Perspective When we observe chemical reactions in the laboratory or in the world around us, we are observing matter at the macroscopic level. Matter is anything that has mass and can be observed. We are so often in contact with matter that we tend to accept our intuitive feel for its existence as an adequate definition. When we study chemistry, however, we need to be aware that some of what we observe in nature is not matter. For example, light is not considered matter because it has no mass. When we take a close look at matter—in this case aluminum—we can see that various questions arise. The behavior of the aluminum in a can is predictable. If the can is tossed into the air, little will happen except that the can will fall to the earth under the force of

4


1.2 The Study of Chemistry


gravity. Aluminum cans and other consumer goods like those shown in Figure 1.2 do not decompose in the air or undergo other chemical reactions. If the aluminum from a soda can is ground into a fine powder and tossed into the air, however, it may ignite—chemically combining with the oxygen in air. It is now believed that the Hindenburg airship burned primarily because it was covered with a paint containing aluminum powder and not because it was filled with hydrogen gas. (You can easily find a summary of the evidence by doing a web search.) One of the most common ways to observe matter is to allow it to change in some way. Two types of changes can be distinguished: physical changes and chemical changes. The substances involved in a physical change do not lose their chemical identities. Physical properties are variables that we can measure without changing the identity of the substance being observed. Mass and density are familiar physical properties. Mass is measured by comparing the object given and some standard, using a balance. Density is a ratio of mass to volume. (This variable is sometimes called mass density). To determine density, both mass and volume must be measured. But these values can be obtained without changing the material, so den- Figure 1.2 ❚ All of the common kitchen items sity is a physical property. Familiar examples of physical properties shown here are made of aluminum. The metal’s light also include color, viscosity, hardness, and temperature. Some other weight, corrosion resistance, and low cost make it a physical properties, which will be defined later, include heat capacity, likely choice for many consumer products. boiling point, melting point, and volatility. Chemical properties are associated with the types of chemical changes that a substance undergoes. For example, some materials burn readily, whereas others do not. Burning in oxygen is a chemical reaction called combustion. Corrosion—the We will discuss corrosion and its degradation of metals in the presence of air and moisture—is another commonly prevention in detail in Chapter 13. observed chemical change. Treating a metal with some other material, such as paint, can often prevent the damage caused by corrosion. Thus an important chemical property of paint is its ability to prevent corrosion. Chemical properties can be determined only by observing how a substance changes its identity in chemical reactions. Both chemical and physical properties of aluminum are important to its utility. A Aluminum is generally found second, structural material is useful only if it can be formed into desired shapes, which requires behind gold, in rankings of metal it to be malleable. Malleability is a measure of a material’s ability to be rolled or ham- malleability. mered into thin sheets, and metals are valuable in part because of their malleability. It is a physical property because the substance remains intact—it is still the same metal, just in a different shape. An aluminum can is formed during its manufacturing process, but its shape can be changed, as you have perhaps done many times when you crushed a can to put it into a recycling bin. Similarly, the chemical properties of aluminum are important. Pure aluminum would be very likely to react with the acids in many popular soft drinks. So aluminum cans are coated inside with a thin layer of polymer—a plastic—to keep the metal from reacting with the contents. This demonstrates how knowing chemical properties can allow product designers to account for and avoid potentially harmful reactions. When we observe chemical reactions macroscopically, we encounter three Two other states of matter are plasmas common states, or phases, of matter: solids, liquids, and gases. At the macroscopic and Bose-Einstein condensates. level, solids are hard and do not change their shapes easily. When a solid is placed in But these do not exist at ordinary a container, it retains its own shape rather than assuming that of the container. Even a temperatures. powdered solid demonstrates this trait because the individual particles still retain their shape, even though the collection of them may take on the shape of the container. Liquids can be distinguished from solids macroscopically because unlike solids, liquids adapt to the shape of the container in which they are held. They may not fill the entire volume, but the portion they do occupy has its shape defined by the container. Finally, gases can be distinguished macroscopically from both liquids and solids primarily because a gas expands to occupy the entire volume of its container. Although many gases are colorless and thus invisible, the observation that a gas fills

5

the available volume is a common experience; when we walk through a large room, we are not concerned that we will hit a pocket with no air. The aluminum that we encounter daily is a solid, but during the refining process, the metal must become molten, or liquid. Handling the molten metal, pouring it into containers, and separating impurities provide both chemical and engineering challenges for those who design aluminum production plants. Often, chemical and physical properties are difficult to distinguish at the macroscopic level. We can assert that boiling water is a physical change, but if you do nothing more than observe that the water in a boiling pot disappears, how do you know if it has undergone a chemical or physical change? To answer this type of question, we need to consider the particles that make up the water, or whatever we observe, and consider what is happening at the microscopic level.

The Microscopic or Particulate Perspective

The word atom comes from the Greek word “atomos” meaning indivisible.

The most fundamental tenet of chemistry is that all matter is composed of atoms and molecules. This is why chemists tend to think of everything as “a chemical” of one sort or another. In many cases, the matter we encounter is a complex mixture of chemicals, and we refer to each individual component as a chemical substance. We will define these terms much more extensively as our study of chemistry develops, but we’ll use basic definitions here. All matter comprises a limited number of “building blocks,” called elements. Often, the elements are associated with the periodic table of elements, shown inside the back cover of this textbook and probably hanging in the room where your chemistry class meets. Atoms are unimaginably small particles that cannot be made any smaller and still behave like a chemical system. When we study matter at levels smaller than an atom, we move into nuclear or elementary particle physics. But atoms are the smallest particles that can exist and retain the chemical identity of whatever element they happen to be. Molecules are groups of atoms held together so that they form a unit whose identity is distinguishably different from the atoms alone. Ultimately, we will see how forces known as “chemical bonds” are responsible for holding the atoms together in these molecules. The particulate perspective provides a more detailed look at the distinction between chemical and physical changes. Because atoms and molecules are far too small to observe directly or to photograph, typically we will use simplified, schematic drawings to depict them in this book. Often, atoms and molecules will be drawn as spheres to depict them and consider their changes. If we consider solids, liquids, and gases, how do they differ at the particulate level? Figure 1.3 provides a very simple but useful illustration. Note that the atoms

To correctly depict the relative densities of a gas and a liquid, much more space would need to be shown between particles in a gas than can be shown in a drawing like Figure 1.3.

Solid

Liquid

Gas

Figure 1.3 ❚ Particulate level views of the solid, liquid, and gas phases of matter. In a solid, the molecules maintain a regular ordered structure, so a sample maintains its size and shape. In a liquid, the molecules remain close to one another, but the ordered array breaks down. At the macroscopic level, this allows the liquid to flow and take on the shape of its container. In the gas phase, the molecules are very widely separated, and move independently of one another. This allows the gas to fill the available volume of the container.

6


Photo: © Cengage Learning/Charles D. Winters

in a solid are packed closely together, and it is depicted as maintaining its shape—here as a block or chunk. The liquid phase also has its constituent particles closely packed, but they are shown filling the bottom of the container rather than maintaining their shape. Finally, the gas is shown with H2O (liquid) H2O (gas) much larger distances between the particles, and the particles themselves move freely through the entire volume of the container. These pictures have been inferred from experiments that have been conducted over many years. Many solids, for example, have well-ordered structures, called crystals, so a particulate representation of solids usually includes this sense of order. How can we distinguish between a chemical and a physical change in this perspective? The difference is much easier to denote at this level, though often it is no more obvious to observe. If a process is a physical change, the atoms or molecules themselves do not change at all. To look at this idea, we turn to a “famous” molecule—water. Many Macroscopic view Microscopic view people who have never studied chemistry can tell you that the chemical formula of water is “H two O.” We depict this Figure 1.4 ❚ The boiling of water is a physical change, in which molecule using different sized spheres; the slightly larger liquid water is converted into a gas. Both the liquid and gas sphere represents oxygen and the smaller spheres represent phases are made up of water molecules; each molecule contains hydrogen. In Figure 1.4, we see that when water boils, the two hydrogen atoms and one oxygen atom. The particulate scale composition of the individual molecules is the same in the insets in this figure emphasize that fact and also show that the liquid phase and the gas phase. Water has not been altered, separation between water molecules is much larger in the gas and this fact is characteristic of a physical change. than in the liquid. Contrast this with Figure 1.5, which depicts a process called electrolysis at the particulate level; electrolysis occurs when water is exposed to an electric current. Notice that the molecules themselves change in this depiction, as water molecules are converted into hydrogen and oxygen molecules. Here, then, we have a chemical change. If we observe these two reactions macroscopically, what would we see and how would we know the difference? In both cases, we would see bubbles forming, only in one case the bubbles will contain water vapor (gas) and in the other they contain hydrogen or oxygen. Despite this similarity, we can make observations at the macroscopic level to distinguish between these two possibilities. Example Problem 1.1 poses an experiment that could be set up to make such an observation. Figure 1.5 ❚ If a suitable electric current is passed through liquid water, a chemical change known as electrolysis occurs. In this process, water molecules are converted into molecules of hydrogen and oxygen gases, as shown in the particulate scale insets in the figure.

Liquid water

Hydrogen gas


Oxygen gas

1.2 The Study of Chemistry

7

E XAM PL E PRO BLEM

1.1

Thomas Holme and Keith Krumnow

Consider the experimental apparatus shown in the photo to the left, in which a candle is suspended above boiling water. This equipment could be used to test a hypothesis about the chemical composition of the gas in the bubbles that rise from boiling water. What would be observed if the bubbles were composed of (a) water, (b) hydrogen, or (c) oxygen?

Strategy This problem asks you to think about what you expect to observe in an experiment and alternatives for different hypotheses. At this stage, you may need to do a little research to answer this question—find out how hydrogen gas behaves chemically in the presence of a flame. We also have to remember some basic facts about fire that we’ve seen in science classes before. To be sustained, fire requires both a fuel and an oxidizer—usually the oxygen in air.

Solution (a) If the bubbles coming out of the liquid contain water, we would expect the flame to diminish in size or be extinguished. Water does not sustain the chemical reaction of combustion (as oxygen does), so if the bubbles are water, the flame should not burn as brightly. (b) You should have been able to find (on the web, for example) that hydrogen tends to burn explosively. If the bubbles coming out of the water were hydrogen gas, we would expect to see the flame ignite the gas with some sort of an explosion. (Hopefully, a small one.) (c) If the bubbles were oxygen, the flame should burn more brightly. The amount of fuel would remain the same, but the bubbles would increase the amount of oxygen present and make the reaction more intense.

Check Your Understanding Work with students in your class or with your instructor to construct this apparatus and see whether or not your observations confirm any of these hypotheses. Draw a picture showing a particulate level explanation for what you observe.

Stan Celestian and Glendale Community College

Symbolic Representation

A sample of bauxite.

8

The third way that chemists perceive their subject is to use symbols to represent the atoms, molecules, and reactions that make up the science. We will wait to introduce this perspective in detail in the next two chapters, but here we point out that you certainly have encountered chemical symbols in your previous studies. The famous “H two O” molecule we have noted is never depicted as we have done here in the quotation marks. Rather, you have seen the symbolic representation of water, H2O. In Chapter 2, we will look at chemical formulas in more detail, and in Chapter 3, we will see how we use them to describe reactions using chemical equations. For now, we simply note that this symbolic level of understanding is very important because it provides a way to discuss some of the most abstract parts of chemistry. We need to think about atoms and molecules, and the symbolic representation provides a convenient way to keep track of these particles we’ll never actually see. These symbols will be one of the key ways that we interact with ideas at the particulate level. How can we use these representations to help us think about aluminum ore or aluminum metal? The macroscopic representation is the most familiar, especially to the engineer. From a practical perspective, the clear differences between unrefined ore and usable aluminum metal are apparent immediately. The principal ore from which aluminum is refined is called bauxite, and bauxite looks pretty much like ordinary


Figure 1.6 ❚ A particulate level representation of aluminum oxide (left) and pure aluminum (right). The gray spheres represent aluminum atoms, and the red spheres represent oxygen.

Aluminum Al

rock. There’s no mistaking that it is different from aluminum metal. At the molecular level, we might focus on the aluminum oxide (also called alumina) in the ore and compare it to aluminum metal, as shown in Figure 1.6. This type of drawing emphasizes the fact that the ore is made up of different types of atoms, whereas only one type of atom is present in the metal. (Note that metals normally contain small amounts of impurities, sometimes introduced intentionally to provide specific, desirable properties. But in this case, we have simplified the illustration by eliminating any impurities.) Finally, Figure 1.6 also shows the symbolic representation for aluminum oxide—its chemical formula. This formula is slightly more complicated than that of water, and we’ll look at this type of symbolism more closely in Chapter 2.

1.3


Aluminum oxide (Alumina) Al2O3

A block of aluminum.

The Science of Chemistry: Observations and Models

Chemistry is an empirical science. In other words, scientists who study chemistry do so by measuring properties of chemical substances and observing chemical reactions. Once observations have been made, models are created to help organize and explain the data. This structure of observations and models provides the backdrop of the science that we’ll explore throughout this book. In some sense, one difference between an engineer and a chemist is that chemists use their intellects and creativity to create models for understanding nature. Usually, the product of intellect and curiosity in engineering is a design that exploits or constrains nature. Ultimately, both fields must begin with the observation of nature.

Observations in Science Observations in chemistry are made in a wide variety of ways for a wide variety of reasons. In some cases, the observations are made because materials with certain properties are needed. For example, containers that hold liquids such as soft drinks need to be strong enough to hold the liquid but light in weight so they don’t increase the cost of transporting the product too much. Before aluminum cans were widely used, steel cans were the containers that society demanded. But steel is relatively heavy, so there was an incentive to find a different packaging material. Scientists and engineers worked together to make observations that confirmed the desirability of aluminum for this use. Observations of nature involve some level of uncertainty in most cases. As an analogy, consider the attendance at a football game. We may be able to count with complete accuracy how many people use tickets to attend a football game, which gives the paid attendance. But that number is not really how many people are there because it does not include members of the press, vendors, and coaching staffs, among others. This example points out one characteristic of making observations: we must be

A single aluminum can has a mass of roughly 14 grams.

1.3 The Science of Chemistry: Observations and Models

9

careful to define what we intend to observe. Carefully defining the measurement to be made, however, does not eliminate all sources of uncertainty. Virtually any scientific measurement must be made more than once to be valid. Because we cannot observe nature with complete certainty, we need to establish the types of uncertainty we encounter in making observations. To do this, two terms are used: accuracy and precision. Although these words may be synonymous in everyday speech, each has its own meaning in scientific or engineering usage. Accuracy indicates how close the observed value is to the “true” value. Precision is the spread in values obtained from the measurement. A precise observation has several measurements close in value. Figure 1.7 illustrates the concepts of accuracy and precision in measurement and shows how the two terms differ. Error in measurement is unavoidable. Again, characteristics of error fall into categories. Random error is fundamental to any measurement. Random error may make the measurement either too high or too low and is associated with the limitations of the equipment with which the measurement is made. Systematic error makes measurements consistently either too high or too low. This type of error is often associated with the existence of some unknown bias in the measurement apparatus. Impurities in metals provide one example of possible error sources. Suppose that an aluminum alloy contains very small amounts of another element, such as silicon. If we attempt to determine the amount of silicon in a sample, we might expect to have significant error in the measurement if the impurity is not uniformly distributed. This will be especially true if we measure only a small sample. Systematic errors are also possible. Consider what would happen if we estimated the density of aluminum using soda cans and did not account for the thin layer of polymer that is applied to the can. Because the density of the polymer coating is likely to be different from that of aluminum, the measurement would be systematically incorrect.

Interpreting Observations


Not all experiments provide direct information about the questions they ask. In many cases, we must infer answers from the data that are obtained. Two types of reasoning that are useful are inductive and deductive reasoning.

(a) Poor precision and poor accuracy

(b) Good precision and poor accuracy

(c) Good precision and good accuracy

Figure 1.7 ❚ The dartboards illustrate the concepts of precision and accuracy and show the difference between the two terms. The bull’s-eye on the target represents the true value of a quantity, and each dart represents a measurement of that quantity. The darts in the left-hand panel are widely scattered and far away from the target; this is neither precise nor accurate. In the center panel, the darts are closely bunched but are not close to the target; this is good precision but poor accuracy. In the right-hand panel, we see good precision and good accuracy because the darts are closely clustered and close to the target. 10


Inductive reasoning begins with a series of specific observations and attempts to generalize to a larger, more universal conclusion. We have asserted that all gases expand to occupy the full volumes of their containers. This universal conclusion was first drawn by inductive reasoning based on observations of many different gases under many different conditions. Because we have not observed every possible gas in every possible container under every possible set of conditions, it could be argued that some counterexample could be found eventually. Deductive reasoning takes two or more statements or assertions and combines them so that a clear and irrefutable conclusion can be drawn. This process can be summarized as an “If A and B, then C” approach that is studied in detail in formal courses in logic. To look at these types of reasoning, consider the developments that led to the affordable refining of aluminum. Even in the 19th century, aluminum ore was available. Metallurgists and other scientists were aware that the combination of aluminum with oxygen in ores is very stable. Observations allowed them to deduce that the chemical bonding forces in the alumina must be strong and would be difficult to overcome. Initial schemes for refining the metal involved heating and adding materials, such as carbon, that would react with oxygen that might be liberated. This method was already in use for refining iron, so inductive reasoning suggested that it might be generalized to other metals, including aluminum. This was a reasonable hypothesis, but it turned out that the methods used in refining iron didn’t work for aluminum. As new discoveries were made about electrochemistry, people realized that these discoveries could lead to a viable scheme for refining aluminum. This is deductive reasoning: an electric current can be used to drive a difficult chemical reaction, and if purifying aluminum involves a difficult chemical reaction, then an electric current must be able to drive the purification of aluminum. Further experiments showed that this was correct, but it was still necessary to make it energetically and economically efficient. Eventually, Charles Hall devised a method that reduced the energy input needed for electrochemical refining of aluminum ore. Later, when this approach was coupled with the relatively low cost of hydroelectricity, the availability of aluminum increased greatly. Ultimately, both deductive and inductive reasoning played vital roles in this technological advance.

Models in Science The number of observations made in the history of science is tremendous. To organize this vast amount of information, scientists create models and theories to make sense of a range of observations. The words “model” and “theory” are sometimes used interchangeably, although some scientists feel that the distinctions between the two are important. Usually, the term model refers to a largely empirical description, such as the fact that gas pressure is proportional to temperature. The word theory, on the other hand, most often refers to an explanation that is grounded in some more fundamental principle or assumption about the behavior of a system. In the connection between gas pressure and temperature, for example, the kinetic theory of gases uses arguments from physics to explain why the molecular nature of gases leads to the observed proportionality. Models are important for a number of reasons. First, they allow us to summarize a large number of observations concisely. Second, they allow us to predict behavior in circumstances that we haven’t previously encountered. Third, they represent examples of creative thinking and problem solving. Finally, constructing and refining models can lead us ultimately to a more fundamental understanding of a problem. Models usually take time to develop. Faced with some initial observations, a creative scientist will try to explain why some pattern exists. The explanation may be qualitative, or it may include a mathematical component to allow quantitative testing or predictions. This insight is considered tentative until additional observations support the hypothesis. The way that science develops, then, involves a cyclic process. Data are obtained, a hypothesis is advanced, and more data are gathered to support or refute

We’ll explore gases and the kinetic theory of gases in Chapter 5.

In popular culture, saying that something is just a theory indicates uncertainty. In science, a theory has significant credibility gained by many observations that confirm its validity.

1.3 The Science of Chemistry: Observations and Models

11

Mass is conserved in ordinary chemical reactions. In nuclear reactions, which we will discuss in Chapter 14, mass and energy can be interconverted.

the hypothesis. A refuted explanation can be discarded or modified to fit the new data and then tested again. Eventually, this process of proposing explanations and testing the implications of those proposals leads to a theory that explains some facet of nature. Models or theories are frequently dynamic, evolving as new information is obtained. Scientists always accept the possibility that a theory can be overturned by new observations that invalidate their explanations of nature. The theories and models we will discuss in this text, however, are supported by a substantial number of experimental measurements and generally have been in place for many years. Although we hear phrases like “it’s just a theory” in casual conversation or even in political discussions of public policy, scientists only use the word to describe a very well-established model. A few theories become so sufficiently refined, well tested, and widely accepted that they come to be known as laws. Because some of these laws have been accepted for hundreds of years, they may appear self-evident. Without serious questioning, for example, we accept the fact that mass should be conserved. Yet the law of conservation of mass did originate in experimental observations of chemical and physical changes carried out by early scientists. As in any theory, a careful scientist is at least implicitly open to the possibility that a principle viewed as a law could be refuted by some unexpected experiment. But, typically, if we encounter a law, we can expect it to summarize nature reliably enough so that we need not be concerned that it will be “repealed.” The process just described is commonly called the scientific method, as we defined it in the “Insight” section of this chapter. The word “method” implies a more structured approach than actually exists in most scientific advancement. Many of the advances of science happen coincidentally, as products of serendipity. The stops and starts that are characteristic of scientific development, however, are guided by the process of hypothesis formation and observation of nature. Skepticism is a key component of this process. Explanations are accepted only after they have held up to the scrutiny of experimental observations. The models that were important in discovering affordable ways to refine aluminum were largely associated with the nature of chemical bonding. We’ll revisit this topic in Chapters 6 and 7 and see how the models of chemical bonding were developed.

1.4

Units can have consequences far more important than confusion over height. NASA attributed the loss of a Mars probe in 1999 to errors associated with the confusion of different units used in the design.

12

Numbers and Measurements in Chemistry

We can observe the world in a variety of ways. For example, we may see a professional basketball player and comment that he is tall, a reasonable statement. If we consider that same person on the basketball court, however, we may want to know how tall he is. In this case, the answer to the question might be “six-ten.” To a basketball fan in the United States, this answer makes sense because that person ascribes additional meaning to the answer, specifically, that the numbers are in feet and inches: 6 feet 10 inches. A fan from another country might not find any meaning in the two numbers, six and ten, because they wouldn’t make sense in metric units. After all, no one is six meters tall. The example of the basketball player offers a good illustration of the difference between qualitative and quantitative information and also points to the need to be specific in the way we communicate information. Science and engineering are regarded as quantitative professions, and usually this reputation is correct. But scientists also look at the world around us in general ways, akin to the assessment that the player is “tall.” Such general or qualitative observations can be crucial in establishing a systematic understanding of nature. As that understanding deepens, though, we usually invoke quantitative or numerical measurements and models. When discussing chemistry, qualitatively or quantitatively, it is always important to communicate our observations and results as clearly as possible. In quantitative observations, this


usually means making sure that we carefully define the terms we use. As we move through this course, you will encounter many examples of “everyday words” that take on much more specific meanings when used in a scientific context. Similarly, when we talk about numerical measurements, we need to be very careful about the ways we use numbers and units.

Units The possible misunderstanding between basketball fans from different countries about the height of a player provides an analogy to an unfortunate reality in studying science and engineering. Science has grown over the centuries, with contributions from a number of cultures and individuals. The legacy of centuries of development is the existence of a large number of units for virtually every basic measurement in science. Energy, for example, has long been important for human civilization and accordingly there are many units of energy. Distance and mass have also been measured for literally millennia. The internationalization of science and engineering led to the establishment of a standard system that provides the needed flexibility to handle a wide array of observations. In this International System of Units (Système International d’Unités, or SI), carefully defined units are combined with a set of prefixes that designate powers of ten. This allows us to report and understand quantities of any size, as illustrated in Figure 1.8. When observations are reported in this system, the base unit designates the type of quantity measured. For example, we can immediately recognize any quantity reported in meters (m) as a distance. Table 1.1 shows the base units for a number of the quantities that chemists may want to measure. These base units, however, are not always conveniently matched to the size of the quantity to be measured. When we consider the sizes of single atoms or molecules, we will see that they are on the order of 0.0000000001 meters. Dealing with so many decimal places invites confusion, so usually it is preferable to choose a unit that is better matched to the scale of the quantity. The standard way to do this is to use prefixes that alter the “size” of any given base unit. A familiar example is the concept of the “kilo” unit. A kilometer is 1000 meters, and a kilobyte, familiar in computer technology, is roughly 1000 bytes. The prefix kiloimplies that there are 1000 of whatever the base unit is. Prefixes such as this exist for numbers large and small, and these prefixes are provided in Table 1.2. The distance of 0.0000000001 meters could be reported as either 0.1 nm or 100 pm. Measuring time Macroscale

1⫻100 m 1m

Height of human

1⫻10–1 m 1 dm

Sheet of paper

Microscale

1⫻10–2 m 1 cm

Wedding ring

1⫻10–3 m 1 mm

Thickness of a CD

1⫻10–4 m 100 µm

1⫻10–5 m 10 µm

Nanoscale

1⫻10–6 m 1 µm

Plant Animal Bacterial cell cell cell

1⫻10–7 m 100 nm

1⫻10–8 m 10 nm

1⫻10–9 m 1 nm

1⫻10–11 m 10 pm 1⫻10–10 m 1⫻10–12 m 100 pm 1 pm

Virus Protein Aspirin Water Atom molecule

Optical microscope Electron microscope Human eye

Because powers of 2 are important in computer science, a kilobyte is actually defined as 1024 bytes.

Specialized techniques are required to observe nanoscale objects.

Scanning tunneling microscope

Figure 1.8 ❚ The typical dimensions of the objects shown extend over many orders of magnitude and help point out the usefulness of the prefixes in the SI system.

1.4 Numbers and Measurements in Chemistry

13

Table

❚

1.1

Base quantities of the SI system of units

The kilogram is considered the base unit for mass, but the names for mass units are derived by affixing prefixes to the gram. A kilogram is 1000 grams.

Table

❚

Property

Unit, with abbreviation

Mass

kilogram, kg

Time

second, s

Distance

meter, m

Electric current

ampere, A

Temperature

kelvin, K

Number of particles

mole, mol

Light intensity

candela, cd

1.2

Prefixes used in the SI system

Factor

Name

Name

Symbol

Y

10

−1

deci

d

zetta

Z

10−2

centi

c

1018

exa

E

10−3

milli

m

1015

peta

P

10−6

micro

μ

1012

tera

T

10−9

nano

n

109

giga

G

10−12

pico

p

106

mega

M

10−15

femto

f

atto

a

24

yotta

1021

10

10

3

Symbol

Factor

−18

kilo

k

10

102

hecto

h

10−21

zepto

z

101

deka

da

10−24

yocto

y

provides another interesting example of units. Laser based experiments can measure the progress of chemical reactions on a timescale of 10−15 s. Therefore, scientific papers that report such experiments use femtoseconds, fs, to report time. Not every quantity can be measured directly in terms of just the seven base units shown in Table 1.1, of course. Some units comprise combinations of these base units and therefore are termed derived units. The SI unit for energy, for example, is the joule (J), and 1 J is defined as 1 kg m2 s−2. In principle, any quantity could be expressed in terms of appropriate combinations of SI base units. But in practice, many other units are so firmly entrenched that they remain in common use. One simple example of this is the measurement of times that are longer than seconds. Minutes, days, and years are used in these circumstances, rather than kiloseconds, and so on. In chemistry labs—and on soda containers—volumes are most often reported in liters (L) or milliliters (mL) rather than the SI unit of cubic meters (m3). Chemists also use a wide variety of units to describe concentration, which measures how much of a particular substance is present

14


in a mixture. Metals often contain minor impurities, and in some cases the units used are simply percentages; other units used include parts per million (ppm) and parts per billion (ppb). The ppm unit tells how many particles of a particular substance are present for every 1,000,000 particles in the sample. In ppb, the sample size is one billion particles. Impurity concentrations on the order of even a few ppm may cause problems in some applications, or they may be added intentionally to impart some desirable property. Later in this chapter, we will take up the issue of converting a measurement from one unit to another, as is frequently required in scientific and engineering calculations. Although many of these SI units have found their way gradually into everyday use, the units for temperature may be the least familiar. You are probably used to seeing temperatures in either degrees Fahrenheit or degrees Celsius, but generally, even in “fully metric” countries, the weather report does not use the Kelvin temperature scale. In general, temperature scales arise from the choice of two standard reference points that can be used to calibrate temperature with the use of a thermometer. The familiar Fahrenheit scale originally chose body temperature as one reference and set it at 100°F. (Accuracy of measurement was clearly not a priority when this temperature scale was proposed.) The second reference point was the coldest temperature that could be achieved by adding salt to ice water, a practice that lowers the melting point of ice. This established 0°F, and the temperature range between the two points was divided into 100 equal units. The scale is now defined by setting the freezing point of water at 32°F and the boiling point of water at 212°F. The Celsius scale was developed in a similar way, but with the freezing point of pure water set at 0°C and the boiling point of water at 100°C. Figure 1.9 shows the relationship between these scales. Conversions between the two scales are given by the following expressions: °F = (1.8 × °C) + 32

(1.1)

°C = (°F − 32) /1.8

(1.2)

Gabriel Daniel Fahrenheit invented the alcohol thermometer in 1709 and the mercury thermometer in 1714.

Scientific uses of temperature require yet another temperature scale. The choice of the kelvin as the standard reflects mathematical convenience more than familiarity. The Kelvin scale is similar to the Celsius scale but draws its utility from the fact that the lowest temperature theoretically possible is zero kelvin. It violates the laws of nature to go below 0 K, as we will see in Chapter 10. The mathematical importance of this definition is that we are assured we will not divide by zero when we use a formula

Fahrenheit Boiling point of water

212°

180° F

Freezing point of water

32°

Celsius 100°

100° C

0°

Kelvin (or absolute) 373

Figure 1.9 ❚ The Fahrenheit, Celsius, and Kelvin temperature scales are compared. The freezing point of water can be expressed as 32°F, 0°C, or 273 K. The boiling point of water is 212°F, 100°C, or 373 K.

100 K

273


15

that has temperature in the denominator of an expression. Conversions between Celsius degrees and kelvins are common in science and are also more straightforward. K = °C + 273.15 °C = K − 273.15

(1.3) (1.4)

Engineers in some disciplines use the Rankine temperature scale (°R), which is an absolute scale whose degrees are the same size as those of the Fahrenheit scale.

Numbers and Significant Figures

Scientific notation provides an advantage here because any digits shown in the number are significant. There are no zeros needed for placing the decimal.

We often encounter very small and very large numbers in chemistry problems. For example, pesticide production in the world exceeds millions of tons, whereas pesticide residues that may harm animals or humans can have masses as small as nanograms. For either type of number, scientific notation is useful. Numbers written using scientific notation factor out all powers of ten and write them separately. Thus the number 54,000 is written as 5.4 × 104. This notation is equivalent to 5.4 × 10,000, which clearly is 54,000. Small numbers can also be written in scientific notation using negative powers of ten because 10−x is identical to 1/10x. The number 0.000042 is 4.2 × 10−5 in scientific notation. When numbers are derived from observations of nature, we need to report them with the correct number of significant figures. Significant figures are used to indicate the amount of information that is reliable when discussing a measurement. “Pure” numbers can be manipulated in a mathematical sense without accounting for how much information is reliable. When we divide the integer 5 by the integer 8, for example, the answer is exactly 0.625, and we would not round this to 0.6. When numbers associated with an observation are used, however, we must be more careful about reporting digits. To understand why, consider whether or not to accept a wager about some measurement. For example, an almanac lists the population of Canada as 33,507,506. Suppose a study concluded that 24% of the people who live in Canada speak French. Based on this information alone, the wager is that 8,041,801 Canadians speak French. Should we accept the bet? What if the wager was that roughly 8.0 million Canadians speak French? Is this a better bet? This scenario shows the importance of significant figures. The figure of 8,041,801 is not believable because we don’t really know if 24% is exactly 24.000000%. It could just as well be 23.95% and the correct answer would then be 8,025,048. Both answers might qualify as “roughly 8 million,” an answer with a more reasonable number of significant figures. This type of reasoning has been formalized into rules for significant figures (digits) in numbers reported from scientific observations. When a measurement is reported numerically, generally we consider that each digit given is known accurately, with one important exception. There are special rules for zero such that it is sometimes significant and sometimes not. When a zero establishes the place for the decimal in the number, it is not significant. Thus the measurement 51,300 m has two zeros that are not significant, as does the measurement 0.043 g. A zero is significant when it is the final digit after a decimal place or when it is between other significant digits. Thus the zeros in both 4.30 mL and 304.2 kg are significant. When numbers are written properly in scientific notation, all of the digits provided are significant. Example Problem 1.2 provides some additional practice in determining the number of significant figures in measurements.

E XAM PL E PRO BLEM

1.2

An alloy contains 2.05% of some impurity. How many significant figures are reported in this value?

Strategy Use the general rule for significant figures: digits reported are significant unless they are zeros whose sole purpose is to position the decimal place.

16


Solution In this case, all digits reported are significant, so there are three significant figures in the number.

Check Your Understanding How many significant figures are reported in each of the following measurements? (a) 0.000403 s, (b) 200,000 g

We also need to account for significant figures in assessing values that we obtain from calculations. The general principle is that a calculated value should be reported with a number of significant figures that is consistent with the data used in the calculation. Our earlier story about the number of French-speaking Canadians provides some insight into this issue as well. The wager that roughly 8.0 million such people live in Canada was the more attractive. The reason: 8.0 million has the same number of significant figures as the value 24%, so it provides a similar quality of information. Three key rules will be required to determine the number of significant figures in the results of calculations. Rule 1: For multiplication and division, the number of significant figures in a result must be the same as the number of significant figures in the factor with the fewest significant figures. When 0.24 kg is multiplied by 4621 m, the result on a calculator reads 1109.04, but if significant figures are correctly reported, the result will be 1100, or 1.1 × 103 kg m. The value 0.24 kg has only two significant figures, and so the result should also have just two significant figures. Rule 2: For addition and subtraction, the rules for significant figures center on the position of the first doubtful digit rather than on the number of significant digits. The result (a sum or difference) should be rounded so that the last digit retained is the first uncertain digit. If the numbers added or subtracted are in scientific notation with the same power of 10, this means that the result must have the same number of digits to the right of the decimal point as in the measurement that has the fewest digits to the right of the decimal point. The number 0.3 m added to 4.882 m yields 5.182 on a calculator, but should be reported as 5.2 m. The first uncertain digit in the values added is the ‘3’ in 0.3 m, so the result should be rounded to one decimal place. In this case, the last allowed digit was “rounded up” because the first nonsignificant digit, 8, was greater than 5. There are two possible conventions for rounding in such cases. In this text, we will round down for numbers 4 and smaller and round up for numbers 5 and larger. Example Problem 1.3 provides some practice in working with significant figures.

E XAM PL E PROB L E M

1. 3

Report the result for the indicated arithmetic operations using the correct number of significant figures. Assume that all values are measurements. (a) 4.30 × 0.31 (b) 4.033 + 88.1 (c) 5.6/1.732 3 104

Strategy Check the rules for significant figures as needed and determine which rules apply. Carry out the calculation and express the result in the correct number of significant figures.

Solution (a) 4.30 × 0.31 = 1.3

(b) 4.033 + 88.1 = 92.1

(c) 5.6/(1.732 × 104) = 3.2 × 10−4

Check Your Understanding Determine the value of the following expressions using the correct number of significant figures: (a) 7.10 m + 9.003 m, (b) 0.004 g × 1.13 g


17

The rules above apply to any numbers that result from most of the measurements that we might make. But in the special case of countable objects, we must also consider one additional rule. Rule 3: When we count discrete objects, the result has no ambiguity. Such measurements use exact numbers, so effectively they have infinite significant figures. Thus if we need to use information such as four quarts in a gallon or two hydrogen atoms in a water molecule, there is no limitation on significant figures. Note that this rule also applies when we work with the various prefixes in the SI system. There are exactly 100 centimeters in a meter, so the factor of 100 would never limit the number of significant figures in a calculation.

1.5

Problem Solving in Chemistry and Engineering

Calculations play a major role in the practice of chemistry and in its application to real world issues and problems. And engineering designs routinely rely on a tremendous number of calculations. The types of exercises we will introduce in this text, though focused on chemistry, will provide practice with techniques that can also be used in engineering applications. To a chemist, the questions associated with aluminum ore require looking into the nature of chemical bonding and how to overcome the stability of strong bonds between aluminum and oxygen. To an engineer, the problems to be addressed in refining the ore might focus on how to deliver enough electricity when and where it is needed. Both disciplines, however, need careful use of quantitative or numerical reasoning.

Using Ratios We encounter and use ratios regularly. We discuss the speed of our cars in miles per hour and buy our fruit at $1.09 a pound. Most of us use such ratios intuitively. If we look at the way we do this, we can deduce a formal set of rules to apply to chemical observations and situations. Let’s say your grocery bill indicates that you paid $4.45 for a 5.0-pound bag of apples. How much did you pay per pound? Although you may not think about it, you obtain the answer by creating an appropriate ratio and completing some simple arithmetic. $4.45 Price = ————— = $0.89 per pound 5.0 pounds The notation “per” means for each of the specified units. Price per pound is the cost of one pound. Miles per hour tell the distance traveled in one hour. This is useful because one is a convenient number in multiplication. We obtain this type of information by creating and simplifying a ratio. With the information we have about the bag of apples, we have another option to create a ratio. We could obtain the number of pounds of apples per dollar. 5.0 pounds $4.45

————— = 1.1 pounds of apples per dollar

For this example, such a manipulation may not be particularly useful—the store probably doesn’t sell 1.1 pound bags of apples. Still, it provides useful insight. When we form the ratio and carry out the indicated arithmetic, the result tells us how much of the numerator is equivalent to one unit of the denominator. So, in general, given two equivalent quantities A = B, we can write either of two ratios: A/B (tells how much A is in one unit of B)

18


and B/A (tells how much B is in one unit of A) Example Problem 1.4 shows how this type of manipulation fits into a problem solving strategy with non-chemistry examples before we look at it in a chemistry context.


1. 4

Shrimp are usually labeled with a “count” that indicates the average number of shrimp per pound. The bigger the shrimp, the smaller the count. Suppose that your supermarket is offering 20-count shrimp for $5.99 per pound. How much should you expect to pay for one dozen shrimp?

Strategy We can determine the average cost per shrimp and then multiply by the number of shrimp needed to find the total price.

Solution $5.99 20 shrimp

————— × 12 shrimp = $3.59

Analyze Your Answer When we calculate a numerical result, it is generally a good idea to pause and see if the answer makes sense. Such a check can often catch errors that might arise from pushing the wrong key on a calculator or even errors in the way we set up the calculation itself. In this case, we found that the cost of 12 shrimp was $3.59, which is a little more than half of the $5.99 price for 20 shrimp. Because 12 is a bit more than half of 20, this seems sensible.

Check Your Understanding A farmer purchases pesticide in 5-gallon drums that cost $23.00 each. If 65 gallons of the pesticide are applied to a field, what was the total cost? How might the situation be more complicated if the amount needed were not a multiple of 5 gallons?

For some problems in chemistry, the units of measurement can be used to determine how to write the appropriate ratio. In the previous example, by considering the units in an algebraic sense, the “shrimp” (with the 20) in the denominator are “canceled” by the same unit (with the 12) in the numerator. $5.99 20 shrimp

————— × 12 shrimp = $3.59

This type of reasoning is called dimensional analysis, or the factor-label method for calculation. In some cases, the dimensions can provide clues about what ratio to define. We will point out this type of reasoning in problems where it may be used.

Ratios in Chemistry Calculations We use ratios in a variety of common calculations in chemistry. One that you will undoubtedly also encounter in engineering is the need to convert between units of different sizes. Example Problem 1.5 introduces this type of manipulation.


1. 5

Visible light is commonly described in terms of its wavelength, which is usually given in units of nanometers. In subsequent calculations, this measurement often needs to be expressed in units of meters. If we are considering orange light of wavelength 615 nm, what is its wavelength in meters?

1.5 Problem Solving in Chemistry and Engineering

19

Strategy We need to establish a ratio that relates nanometers and meters. Then we can use that ratio to determine the answer.

Solution 1 m = 1 × 109 nm We can write this as a ratio. Because we want to convert from nm to m, we’ll need m in the numerator and nm in the denominator. 1m 10 nm

———— 9

Then we just complete the calculation: 1 m = 6.15 × 10–7 m 615 nm × ———— 109 nm

Analyze Your Answer Writing units on all quantities involved helps ensure that we use the correct ratio. Even if we aren’t familiar with the typical wavelengths of light, we can still check that the answer makes sense. Because nanometers are much smaller than meters, the fact that we got a much smaller number for our answer seems correct. Note that conversions between SI units are between exact numbers when considering significant figures.

Check Your Understanding Large amounts of electrical power, measured in watts (W), are used to produce aluminum from ore. If a production plant uses 4.3 GW in a certain period of time, how many W are used?

Another common type of problem that uses ratios occurs when two quantities are related to each other either by some chemical or physical property. One such property is mass density. Mass density is defined as the mass of a substance per unit volume. This definition is itself a ratio that allows us to convert between mass and volume, as shown in Example Problem 1.6.

E XAM PL E PRO BLEM

1.6

The mass density of water at 25°C is 0.997 g per mL. A child’s swimming pool holds 346 L of water at this temperature. What mass of water is in the pool?

Strategy We want to use the density as a ratio, but first we must express the volume in appropriate units. We can convert the volume of the pool from liters to milliliters and then use the density as given.

Solution 1 L = 1000 mL, so 1000 mL/1 L provides a conversion from L to mL. The given density is 0.997 g/1 mL, and this provides the connection between volume and mass. 0.997 g 1000 mL 346 L × ————— × ———— = 3.45 × 105 g 1L 1 mL

Analyze Your Answer It may be difficult to have an intuitive feel for the size of an answer. But 346 L would be nearly 100 gallons of water, which would be fairly heavy. A pound is equivalent to 454 g, so the answer we found (3.45 × 105 g) is more than 700 pounds, and this is at least plausible. If we had inverted the ratio used to convert from mL to L, for example, we would have gotten an answer of just 345 g, or less than a pound. Clearly, that would be too small for even a very tiny pool, so we would have had a chance to find and correct the error.

20


Discussion Here we have written our solution as a single two-step process in which we first converted the volume from liters to milliliters and then used the density to find the mass. Some students feel more comfortable breaking these steps down into separate calculations; we could have found the volume explicitly as 346,000 mL and then multiplied that by the density. As long as the operations are carried out correctly, the result will be the same. If you do choose to break calculations down into smaller pieces, though, be careful not to round your answers until you reach the final result.

Check Your Understanding A liquid pesticide has a density of 1.67 g per mL. What is the mass of liquid in a full 20.0-L container?

Calculations in chemistry use ratios very often. Example Problem 1.6 also shows how they are used in combinations. Learning how to carry out these sequential manipulations will be a key part of solving many of the problems we encounter in chemistry. One guide that often helps to determine how to construct the appropriate ratio is to note the units used. For example, the equality between 1000 mL and 1 L leads to two ways to express the ratio, 1000 mL/1 L or 1 L/1000 mL. Which ratio did we need? Because we started with liters, we required the former ratio, with 1 L in the denominator. We will often include discussions about the ratios we form in the “strategy” part of the example problems. Mass density is more than a convenient way to define a ratio for calculations. It is also an important characteristic of materials. Among the reasons that aluminum is a popular structural material is that its relatively low density is often important. When a design calls for a strong but lightweight material, aluminum is a good candidate. We’ll look at this aspect further in the closing insight in Section 1.6 when we look at various materials used in bicycle frames.

Mass density also is important in determining buoyancy because less dense objects float on more dense ones.

Conceptual Chemistry Problems Although numerical calculations will always be an important component of chemistry, they are only part of the field. Sometimes, to ensure that the concepts involved in chemistry are understood, we’ll also work problems that focus on the particulate representation and other concepts. Often, the strategies employed in this type of question are different from those we have just described. When we try to visualize a chemical concept using drawings, we will start from the conventions that were introduced for the particulate perspective of chemistry in Section 1.2. In this type of problem, we could be asked, for example, to draw a diagram that depicts what happens to the molecules when steam condenses into liquid water. Example Problem 1.7 illustrates this type of thinking.


1. 7

Dry ice is solid carbon dioxide; it is called “dry” in large part because it goes directly from a solid to a gas without becoming a liquid under ordinary conditions. Draw a picture that shows what the carbon dioxide molecules might look like as a solid and as a gas.

Strategy To answer conceptual questions such as this, we must depict the differences between the two phases. Our drawing can be schematic, but it must imply the differences clearly. In this case, we know that a solid will have closely packed molecules, whereas a gas will exhibit significant space between particles.

Solution Our picture shows several things. First, we have included some information about the chemical composition of the molecules by using two different colors for the spheres that represent the atoms. Second, the solid is indicated at the bottom and is distinguishable because the molecules are shown closely packed together in


21

an orderly array. Finally, the gas is depicted by only a few molecules, and they are widely spaced.

Discussion Obviously, the solutions to conceptual questions like this are less exact than numerical answers. So your drawing might look somewhat different, yet still be “correct.” Neither the gas nor the solid needs to look exactly like the picture we have drawn, but the drawing itself should impart the essential concepts we need to understand when thinking about the particulate perspective of this process.

Check Your Understanding Draw a picture that shows a molecular scale view of steam condensing into liquid water.

We will continue to use conceptual problems throughout the text to help provide practice in thinking about the behavior of molecules and atoms.

Visualization in Chemistry

The chemical name for caustic soda is sodium hydroxide, and its formula is NaOH. Lime is calcium oxide, CaO.

One area where conceptual understanding in chemistry is different from that in most engineering pursuits is the way we visualize systems. Chemistry provides multiple simultaneous ways to view problems, including the completely abstract perspective of atoms and molecules that will never be observed directly. Methods to visualize this level of chemistry provide an important tool in the way that chemistry is taught and learned. We can explore this idea by thinking about the refining of aluminum. The most common aluminum ore is bauxite, which consists of aluminum oxides and rock. The first step in producing aluminum metal is separating the aluminum oxides from the remainder of the rock. As shown in Figure 1.10, this actually entails Mining

Liquor for recycling

Pregnant liquor

Bauxite storage

Flash tank Digestion Filter press Blow-off tank

Precipitated aluminum hydroxide

Digestion vessel Caustic soda and lime Steam

Slurry mixing Clarification

Mud lake Based on a drawing by Robert J. Lancashire, University of West Indies

Conveyor

Precipitation

Mud settler

To Red-Mud Lake

Kiln Calcination and shipping

Figure 1.10 ❚ Several steps in the processing of bauxite are represented in this diagram.

22


Ball mill

Alumina (Al2O3)

Silica (SiO2)

Aluminates, dissolved in liquid phase

Unreacted silica

Figure 1.11 ❚ Bauxite is treated with caustic soda and lime in a process called digestion. The alumina in the ore reacts and dissolves. This step separates the aluminum from the rest of the ore, which contains a variety of other minerals.

Treat with lime or caustic soda

several steps; the first is digestion of the ore. Process engineers design “digesters” in which crushed ore, caustic soda, and lime are mixed at high temperatures to create a slurry. If we look at Figure 1.11, we can visualize this process at the microscopic level. The molecular level pictures show alumina as atoms of aluminum and oxygen, and the rock is shown as silicon and oxygen (or silica). This is a simplified depiction based on the fact that much of the earth’s crust contains a large percentage of these elements. (Bauxite ore also contains significant amounts of iron oxide and other minerals, but these behave much like silica in the digestion process.) Using this level of visualization to consider the reactions in the digester, we see that alumina reacts with the caustic soda and lime, but silica does not. In this way, we can look at a large-scale industrial process while thinking in the microscopic perspective. The next step in aluminum refining is smelting. At this stage, a chemical reaction is induced where the aluminates, now dissolved in a material called cryolite (Na3AlF6), lose oxygen atoms to a carbon rod, forming relatively pure aluminum and carbon dioxide. This portion of the process is shown in Figure 1.12. Although this molecular Carbon rod, immersed in cryolite

Note that the conceptual drawings in Figures 1.11 and 1.12 are schematic. They do not attempt to impart molecular detail but rather aim to provide a sense of the particles involved in the chemical processes.

Figure 1.12 ❚ These particulate Carbon dioxide gas carries away oxygen atoms.

level illustrations provide a simplified view of the atomic-scale process involved in smelting aluminum.

Molten cryolite, containing dissolved aluminum oxides

Aluminum falls to bottom of tank and is drawn off for further purification.


23

A metric ton is equal to 1000 kg, or roughly 2200 pounds.

level visualization again shows how the chemical reaction rearranges atoms at the particulate level, it is worth remembering that this process is carried out on a massive scale: the U.S. aluminum industry produces about 2.6 million metric tons of aluminum per year. The visualization techniques introduced in this section will help us to develop a particulate level understanding of many concepts we encounter throughout the text. At this point, let’s take some time to look at one application where the properties of aluminum metal make it a popular material choice for engineering a consumer product—a bicycle.

INSIGHT INTO

1.6

Measurements such as elastic modulus or yield strength depend on temperature. For most engineering designs, however, this dependence is more critical in plastics than in metals.

Material Selection and Bicycle Frames

For the average rider who doesn’t wish to spend a great deal of money on a bike, the frame is likely to be some alloy of steel. Using other materials for the frame can improve certain aspects of performance but also increases the cost of the bicycle. The consumer’s choices are confronted first by the engineer who designs the frame. Among the properties that an engineer must consider in choosing a material for a frame are strength, density (which affects the weight of the frame), and stiffness. Although we are already familiar with density, talking about the strength and stiffness of materials requires some additional vocabulary. Stiffness is related to a property called the elastic modulus of the material. The elastic modulus measures the amount of stretch or compression a material experiences when it is stressed. Something with a high modulus will stretch very little even when experiencing large forces, and a bicycle frame composed from such a material would seem stiff. We will see this concept again when we discuss chemical bonding because the ability of a material to stretch is related to how closely and strongly the atoms in it are connected to each other. The strength of a material is formally measured by a property called the yield strength. The yield strength measures the amount of force required to produce a specified deformation of the material. A stronger material can withstand greater forces before it deforms and so has higher yield strength. Finally, because there are benefits from having a lighter frame, materials with low densities are also desirable. The three most common metals for bicycle frames are aluminum, steel, and titanium. Table 1.3 shows typical values for the elastic modulus, yield strength, and density of these materials. Looking at this table, we can see that the physical properties of these materials afford trade-offs in types of performance in the frame. Although aluminum is lightweight, it is neither as strong nor as stiff as steel or titanium. Steel is the strongest, stiffest, and heaviest of the materials. It’s important to note the range of yield strengths in this table because the details of manipulating the strength of aluminum or steel depend on both the chemical composition and the manner in which the material is treated in fabricating the frame.

Table

❚

1.3

Elastic modulus, yield strength, and density of some materials used in bicycle frames

Elastic Modulus (psi)

Yield Strength Range (psi)

Density (g/cm3)

Aluminum

10.0 × 106

5.0 × 103 − 6.0 × 104

2.699

Steel

30.0 × 106

4.5 × 104 − 1.6 × 105

7.87

Titanium

16.0 × 106

4.0 × 104 − 1.2 × 105

4.507

Material

24


Of course, the engineer has more than the choice of material available in making a design for a bicycle frame. The size and connectivity of the tubing can also be adjusted, for example. Aluminum frames generally feature tubing with a much larger diameter than that in steel or titanium frames. These larger diameter tubes give the frame itself a much stiffer feel—so much so that many cyclists find aluminum frames too stiff to give a comfortable ride. Many modern frames also use tubing that is not round but rather is oval to withstand specific types of stress commonly encountered in cycling. Some very expensive frames for competitive cycling are made from more exotic materials. New aluminum alloys containing trace amounts of scandium, for example, can achieve lower frame weights without the large tubing diameters that cause the jarring ride for which aluminum is usually known. The extremely aerodynamic bikes preferred for high-speed time trial racing are most often made from carbon fiber composites that are molded easily into exotic shapes that reduce air resistance. Chemistry and chemists have played an essential role in helping engineers develop and exploit these new materials. And the intersection between chemistry and engineering illustrated in this chapter for bicycles and soda cans is also apparent in many other familiar products. Throughout this text we’ll look at places where the chemistry we learn has an impact on engineering designs. At this point, we know enough to say that if a lightweight bicycle frame is your desire, you may want to take advantage of aluminum and its low mass density. In Chapter 2, we take the next step and look at atoms and molecules in more detail and simultaneously introduce polymers—another important class of engineering materials.

FO C U S ON

PROB L E M S O LVI NG

Engineering students often wonder whether their college chemistry courses will prove useful later in their careers. Through the opening and closing “Insight” sections of the chapters of this book, we’ll make connections between the content of the chemistry course and places where that content is relevant in engineering applications. But this type of connection isn’t the only place where things learned in chemistry courses help future engineers. Engineers solve problems. And the exposure to various problem solving techniques in courses like this will help you to develop and diversify your skills. We will emphasize this connection after each chapter with a special section like this one. In these “Focus on Problem Solving” sections, we will look at chemistry problems related to the chapter just presented. The difference will be that the “correct” answer will not be a number but rather the identification of an appropriate strategy for approaching the problem.

Question Describe how you would determine which has a greater mass, a sphere of iron shot with a radius of 4.00 mm or a cube of nickel with an edge length of 4.00 mm. Which formulas are needed, and what other information would have to be looked up?

Strategy Like many real world problems, this question does not have enough information given to provide an answer. To devise a strategy, we have to think about what it takes to compare two materials—in this case iron and nickel. Moreover, we are asked to compare masses of the two materials, so that provides a clue. The dimensional information given, along with formulas that are familiar or can be looked up, provides information about volume. Mass and volume were related in this chapter by density, and density values for common materials can be found easily.

Solution To answer this problem, first we need to look up the densities of both iron and nickel. Then, we need to use the data provided to determine the volume of each

Focus on Problem Solving

25

sample. For a cube, we can calculate the volume by taking the given measure for the side (s) and cubing it (s 3). The volume of a sphere is given by (4/3)pr 3. Multiplying the density by volume provides the mass, so with the data we looked up for density and the calculated volumes we could answer the question.

S U MMA RY Chemistry is the science of matter, and since all engineering designs involve matter, the links between chemistry and engineering are many. We began to explore the role of materials in engineering by considering aluminum. By applying some very simple chemical concepts, we can begin to understand the transformation of aluminum from a precious metal to a common and inexpensive material. One common trait of an experienced chemist is the ability to consider a given situation from a number of perspectives. Both the physical and chemical properties of substances can be considered at the macroscopic or microscopic (particulate) level depending on the nature of the question or problem being considered. In addition, chemists often use symbolic representations to describe what is happening in chemical systems. Becoming comfortable with these different perspectives can give students an edge in understanding many chemistry problems. Chemistry is an empirical science. It relies on experimental observations to develop an understanding of matter. The path from observation toward an understanding of the universe typically involves several steps, relying on inductive or deductive reasoning, or both. By applying reasoning skills to our observations, we construct models for understanding chemical phenomena. Then these models are refined and adapted over time. The ideas that we will explore in this text involve many models or theories that have been developed through the scientific method.

Many of the observations of nature that are used to develop theories and models need to be quantitative, that is, they must assess what is being observed with some level of numerical detail. The need for numerical observations throughout the development of chemistry (and other sciences) has given rise to systematic ways to communicate this information. A number alone is not sufficient to impart all the meaning of a measurement; experimental observations are expected to include units of measurement. An important skill in the study of both chemistry and engineering is the ability to manipulate numerical information, including the units attached to that information. The use of ratios to convert between a measurement in one unit and desired information in another related unit represents a core skill for problem solving in chemistry and engineering. The method of dimensional analysis, sometimes called the factor-label method, provides one common way to carry out these transformations. We will build on these fundamental ideas as we proceed in the study of chemistry. The ability to look at problems from several perspectives, extract and manipulate numerical information, and ultimately gain a broad understanding of the chemical principles that underlie the behavior of the universe will provide an interesting challenge as we survey the connections between chemistry and engineering in this book.

KEY TERMS accuracy (1.3)

laws (1.3)

precision (1.3)

atoms (1.2)

liquids (1.2)

random error (1.3)

chemical properties (1.2)

macroscopic perspective (1.2)

scientific method (1.1)

combustion (1.2)

malleability (1.2)

scientific models (1.3)

conceptual chemistry problems (1.5)

mass density (1.2)

scientific notation (1.4)

deductive reasoning (1.3)

matter (1.2)

significant figures (1.4)

density (1.2)

microscopic perspective (1.2)

solids (1.2)

dimensional analysis (1.5)

molecules (1.2)

symbolic perspective (1.2)

elastic modulus (1.6)

particulate perspective (1.2)

systematic error (1.3)

elements (1.2)

parts per billion (ppb) (1.4)

temperature scales (1.4)

factor-label method (1.5)

parts per million (ppm) (1.4)

units (1.4)

gases (1.2)

phases of matter (1.2)

visualization in chemistry (1.5)

inductive reasoning (1.3)

physical properties (1.2)

yield strength (1.6)

26


P ROB L E M S AND E XE RCI S ES ■

denotes problems assignable in OWL

INSIGHT INTO Aluminum 1.1 Use the web to determine the mass of a steel beverage can from the 1970s and the mass of a modern aluminum can. How much more would a 12-pack of soda weigh in steel cans? 1.2 Which properties of aluminum might concern you if you had to use the aluminum tableware that Napoleon employed to impress his guests? 1.3 Where does the scientific method “start”? What is the first step? 1.4 Use the web to determine the amount of aluminum used in the United States in a single year. What is the primary use for this material? 1.5 Use the web to find current prices offered for aluminum for recycling. Is there variation in the price based on where in the United States the aluminum is returned? 1.6 Use the web to determine the differences in the amounts of aluminum recycled in states where there are deposits on aluminum cans versus states where recycling is voluntary. What is the most reliable way to estimate this value? What uncertainty is there in the estimate? The Study of Chemistry 1.7 When we make observations in the laboratory, which perspective of chemistry are we normally using? 1.8 Which of the following items are matter and which are not? (a) a flashlight, (b) sunlight, (c) an echo, (d) air at sea level, and (e) air at the top of Mount Everest. 1.9 Which macroscopic characteristics differentiate solids, liquids, and gases? (List as many as possible.) 1.10 How can a liquid be distinguished from a fine powder? What type of experiment or observation might be undertaken? 1.11 Some farmers use ammonia, NH 3, as a fertilizer. This ammonia is stored in liquid form. Use the particulate perspective to show the transition from liquid ammonia to gaseous ammonia. 1.12 Do the terms element and atom mean the same thing? If not, how do they differ? 1.13

■ Label each of the following as either a physical process or a chemical process: (a) rusting of an iron bridge, (b) melting of ice, (c) burning of a wooden stick, (d) digestion of a baked potato, (e) dissolving of sugar in water

1.14 Why do physical properties play a role in chemistry if they do not involve any chemical changes? 1.15 Physical properties may change because of a chemical change. For example, the color of an egg white changes from clear to white because of a chemical change when it is cooked. Think of another common situation when a chemical change also leads to a physical change. 1.16

■

■ Which

part of the following descriptions of a compound or element refers to its physical properties and which to its chemical properties?

Problems assignable in OWL

(a) Calcium carbonate is a white solid with a density of 2.71 g/cm3. It reacts readily with an acid to produce gaseous carbon dioxide. (b) Gray powdered zinc metal reacts with purple iodine to give a white compound. 1.17 Use a molecular level description to explain why gases are less dense than liquids or solids. 1.18 All molecules attract each other to some extent, and the attraction decreases as the distance between particles increases. Based on this idea, which state of matter would you expect has the strongest interactions between particles: solids, liquids, or gases? Observations and Models 1.19 We used the example of attendance at a football game to emphasize the nature of observations. Describe another example where deciding how to count subjects of interest could affect the observation. 1.20 Complete the following statement: Data that have a large random error but otherwise fall in a narrow range are (a) accurate, (b) precise, or (c) neither. 1.21 Complete the following statement: Data that have a large systematic error can still be (a) accurate, (b) precise, or (c) neither. 1.22 Two golfers are practicing shots around a putting green. Each golfer takes 20 shots. Golfer 1 has 7 shots within 1 meter of the hole, and the other 13 shots are scattered around the green. Golfer 2 has 17 shots that go into a small sand trap near the green and 3 just on the green near the trap. Which golfer is more precise? Which is more accurate? 1.23 Use your own words to explain the difference between deductive and inductive reasoning. 1.24 Suppose that you are waiting at a corner for a bus. Three different routes pass this particular corner. You see buses pass by from the two routes that you are not interested in taking. When you say to yourself, “My bus must be next,” what type of reasoning (deductive or inductive) are you using? Explain your answer. 1.25 When a scientist looks at an experiment and then predicts the results of other related experiments, which type of reasoning is she using? Explain your answer. 1.26 What is the difference between a hypothesis and a question? 1.27 Should the words theory and model be used interchangeably in the context of science? Defend your answer using information found in a web search. 1.28 What is a law of nature? Are all scientific laws examples of laws of nature? Numbers and Measurements 1.29 Describe a miscommunication that can arise because units are not included as part of the information. 1.30 What is the difference between a qualitative and a quantitative measurement?

Problems and Exercises

27

1.31 Identify which of the following units are base units in the SI system: grams, meters, joules, liters, amperes.

(c) x = (0.0061020)(2.0092)(1200.00) (0.0034)2 + 4(1.000)(6.3 × 10−4 ) √‾‾‾‾‾‾‾‾‾‾‾‾‾‾‾ (d) x = 0.0034 + ———————————————— (2)(1.000)

1.32 What is a “derived” unit? 1.33 Rank the following prefi xes in order of increasing size of the number they represent: centi-, giga-, nano-, and kilo-. 1.34 The largest computers now include disk storage space measured in petabytes. How many bytes are in a petabyte? (Recall that in computer terminology, the prefix is only “close” to the value it designates in the metric system.) 1.35 Historically, some unit differences reflected the belief that the quantity measured was different when it was later revealed to be a single entity. Use the web to look up the origins of the energy units erg and calorie, and describe how they represent an example of this type of historical development. 1.36 Use the web to determine how the Btu was initially established. For the engineering applications where this unit is still used today, why is it a sensible unit? 1.37 How many micrograms are equal to one gram? 1.38 Convert the value 0.120 ppb into ppm. 1.39 How was the Fahrenheit temperature scale calibrated? Describe how this calibration process reflects the measurement errors that were evident when the temperature scale was devised. 1.40

Superconductors are materials that have no resistance to the flow of electricity, and they hold great promise in many engineering applications. But to date superconductivity has only been observed under cryogenic conditions. The highest temperature at which superconductivity has been observed is 138 K. Convert this temperature to both °C and °F.

■

1.41 Express each of the following temperatures in kelvins: (a) −10.°C, (b) 0.00°C, (c) 280.°C, (d) 1.4 × 103°C 1.42

■ Express (a) 275°C in K, (b) 25.55 K in °C, (c) −47.0°C in °F, and (d) 100.0°F in K

1.43 Express each of the following numbers in scientific notation: (a) 62.13, (b) 0.000414, (c) 0.0000051, (d) 871,000,000, (e) 9100 1.44

1.48

1.49 A student finds that the mass of an object is 4.131 g and its volume is 7.1 mL. What density should be reported in g/mL? 1.50 Measurements indicate that 23.6% of the residents of a city with a population of 531,314 are college graduates. Considering significant figures, how many college graduates are estimated to reside in this city? 1.51 A student weighs 10 quarters and finds that their total mass is 56.63 grams. What should she report as the average mass of a quarter based on her data? 1.52

Problem Solving in Chemistry and Engineering 1.53 A package of eight apples has a mass of 1.00 kg. What is the average mass of one apple in grams? 1.54 If a 1.00-kg bag containing eight apples costs $1.48, how much does one apple cost? What mass of apples costs $1.00? 1.55 A person measures 173 cm in height. What is this height in meters? feet and inches? 1.56 The distance between two atoms in a molecule is 148 pm. What is this distance in meters? 1.57 Carry out the following unit conversions: (a) 3.47 × 10−6 g to μg, (b) 2.73 × 10−4 L to mL, (c) 725 ns to s, (d) 1.3 m to km ■ Carry out each of the following conversions: (a) 25.5 m to km, (b) 36.3 km to m, (c) 487 kg to g, (d) 1.32 L to mL, (e) 55.9 dL to L, (f ) 6251 L to cm3

1.59

■

1.60

■ If a vehicle is traveling 92 m/s, what is its velocity in miles per hour? (0.62 miles = 1.00 km)

Perform these calculations and express the result with the proper number of significant figures.

1.61

■ A load of asphalt weights 254 lbs. and occupies a volume of 220.0 L. What is the density of this asphalt in g/L?

(a) (4.850 g − 2.34 g)/1.3 mL

1.62

■

1.63

■

■ How

many significant figures are there in each of the following? (a) 0.136 m, (b) 0.0001050 g, (c) 2.700 × 10−3 nm, (d) 6 × 10−4 L, (e) 56003 cm3 ■ How many significant figures are present in these measured quantities? (a) 1374 kg, (b) 0.00348 s, (c) 5.619 mm, (d) 2.475 × 10−3 cm, (e) 33.1 mL

1.46

■

(b) V = pr , where r = 4.112 cm 3

(c) (4.66 × 10−3) × 4.666 (d) 0.003400/65.2 Calculate the following to the correct number of significant figures. Assume that all these numbers are measurements. ■

(a) x = 17.2 + 65.18 − 2.4 (b) x =13.0217/17.10

28

■ A rock is placed on a balance and its mass is determined as 12.1 g. When the rock is then placed in a graduated cylinder that originally contains 11.3 mL of water, the new volume is roughly 17 mL. How should the density of the rock be reported?

1.58

1.45

1.47

■ In an attempt to determine the velocity of a person on a bicycle, an observer uses a stopwatch and finds the length of time it takes to cover 25 “squares” on a sidewalk. The bicycle takes 4.82 seconds to travel this far. A measurement of one of the squares shows that it is 1.13 m long. What velocity, in m/s, should the observer report?


Convert 22.3 mL to (a) liters, (b) cubic inches, and (c) quarts.

One square mile contains exactly 640 acres. How many square meters are in one acre? A sample of crude oil has a density of 0.87 g/mL. What volume in liters does a 3.6-kg sample of this oil occupy?

1.64 Mercury has a density of 13.6 g/mL. What is the mass of 4.72 L of mercury? 1.65

The area of the 48 contiguous states is 3.02 × 106 mi2. Assume that these states are completely flat (no mountains and no valleys). What volume of water, in liters, would cover these states with a rainfall of two inches? ■

■


■ The

dimensions of aluminum foil in a box for sale in supermarkets are 66 2/3 yards by 12 inches. The mass of the foil is 0.83 kg. If its density is 2.70 g/cm3, then what is the thickness of the foil in inches?

1.79 Use the web to research the differences in the design of steel-framed bicycles versus aluminum-framed bicycles. Write a short paragraph that details the similarities and differences you discover.

Titanium is used in airplane bodies because it is strong and light. It has a density of 4.55 g/cm3. If a cylinder of titanium is 7.75 cm long and has a mass of 153.2 g, calculate the diameter of the cylinder. (V = pr 2h, where V is the volume of the cylinder, r is its radius, and h is the height.)

1.80 Use the web to research the elastic modulus and yield strength of carbon fiber composites. How do these materials compare to aluminum, steel, and titanium?

1.66

1.67

■

1.68

■ Wire is often sold in pound spools according to the wire gauge number. That number refers to the diameter of the wire. How many meters are in a 10-lb spool of 12-gauge aluminum wire? A 12-gauge wire has a diameter of 0.0808 in. and aluminum has a density of 2.70 g/cm3. (V = pr 2l )

1.69

■ An

industrial engineer is designing a process to manufacture bullets. The mass of each bullet must be within 0.25% of 150 grains. What range of bullet masses, in mg, will meet this tolerance? 1 gr = 64.79891 mg.

1.70

■ An

engineer is working with archaeologists to create a realistic Roman village in a museum. The plan for a balance in a marketplace calls for 100 granite stones, each weighing 10 denarium. (The denarium was a Roman unit of mass: 1 denarium = 3.396 g.) The manufacturing process for making the stones will remove 20% of the material. If the granite to be used has a density of 2.75 g/cm3, what is the minimum volume of granite that the engineer should order?

1.71 Draw a molecular scale picture to show how a crystal differs from a liquid. 1.72 Draw a molecular scale picture that distinguishes between alumina and silica. Is your picture structurally accurate or schematic? 1.73

■ On average, Earth’s crust contains about 8.1% aluminum by mass. If a standard 12-ounce soft drink can contains approximately 15.0 g of aluminum, how many cans could be made from one ton of the Earth’s crust?

1.74

■ As computer processor speeds increase, it is necessary for engineers to increase the number of circuit elements packed into a given area. Individual circuit elements are often connected using very small copper “wires” deposited directly onto the surface of the chip. In some current generation processors, these copper interconnects are about 65 nm wide. What mass of copper would be in a 1-mm length of such an interconnect, assuming a square cross section? The density of copper is 8.96 g/cm3.

INSIGHT INTO Material Selection and Bicycle Frames 1.75 Suppose that a new material has been devised with an elastic modulus of 22.0 × 106 psi for a bicycle frame. Is this bike frame likely to be more or less stiff than an aluminum frame? 1.76 Rank aluminum, steel, and titanium in order of increasing stiffness. 1.77 Compare the strengths of aluminum, steel, and titanium. If high strength were needed for a particular design, would aluminum be a good choice? 1.78 Aluminum is not as strong as steel. What other factor should be considered when comparing the desirability of aluminum versus steel if strength is important for a design?

■


1.81 Use the web to research the relative cost of aluminum, steel, and titanium frames for bicycles. Speculate about how much of the relative cost is due to the costs of the materials themselves.

FOCUS ON PROBLEM SOLVING EXERCISES 1.82 A student was given two metal cubes that looked similar. One was 1.05 cm on an edge and had a mass of 14.32 grams; the other was 2.66 cm on a side and had a mass of 215.3 grams. How can the student determine if these two cubes of metal are the same material using only the data given? 1.83 Battery acid has a density of 1.285 g/mL and contains 38.0% sulfuric acid by mass. Describe how you would determine the mass of pure sulfuric acid in a car battery, noting which item(s) you would have to measure or look up. 1.84 Unfermented grape juice used to make wine is called a “must.” The sugar content of the must determines whether the wine will be dry or sweet. The sugar content is found by measuring the density of the must. If the density is lower than 1.070 g/mL, then sugar syrup is added until the density reaches 1.075 g/mL. Suppose that you have a sample taken from a must whose mass is 47.28 g and whose volume is 44.60 mL. Describe how you would determine whether or not sugar syrup needs to be added and if so, how would you estimate how much sugar syrup to add? 1.85 A solution of ethanol in water has a volume of 54.2 mL and a mass of 49.6 g. What information would you need to look up and how would you determine the percentage of ethanol in this solution? 1.86 Legend has it that Archimedes, a famous scientist of Ancient Greece, was once commanded by the king to determine if a crown he received was pure gold or a gold−silver alloy. He was not allowed, however, to damage the crown (by slicing off a piece, for example). If you were assigned this same task, what would you need to know about both gold and silver, and how would you make a measurement that would tell you if the crown was pure gold? 1.87 Imagine that you place a cork measuring 1.30 cm × 4.50 cm × 3.00 cm in a pan of water. On top of this cork, you place a small cube of lead measuring 1.15 cm on a side. Describe how you would determine if the combination of the cork and lead cube will still float in the water. Note any information you would need to look up to answer the question. 1.88 A calibrated flask was filled to the 25.00-mL mark with ethyl alcohol and it was found to have a mass of 19.7325 g. In a second experiment, 25.0920 g of metal beads were put into the container and the flask was again filled to the 25.00-mL mark. The total mass of the metal plus the alcohol was 43.0725 g. Describe how to determine the density of the metal sample.


29

2

Atoms and Molecules

OUTLINE 2.1

INSIGHT INTO Polymers

2.2

Atomic Structure and Mass

2.3

Ions

2.4

Compounds and Chemical Bonds

2.5

The Periodic Table

2.6

Inorganic and Organic Chemistry

2.7

Chemical Nomenclature

2.8

INSIGHT INTO Polyethylene

The plastics in consumer and industrial products are examples of polymers. The injection molded samples shown here are used in rigid packaging and durable goods. The Dow Chemical Company

A A variety of techniques known as scanning probe microscopy can produce images of single atoms under some conditions.

toms and molecules are the building blocks of chemistry. You’ve probably been hearing this since middle school, so the existence of atoms is not something that you are likely to question or challenge. Chances are that you rarely think about atoms or molecules, however, when you come across items in your day-to-day life. When chemists want to understand some aspect of the world around them, though, they focus their attention at the level of atoms and molecules. So an important part of studying chemistry is learning how to interpret nature by thinking about what atoms and molecules are doing. Certainly, it can take some time to get comfortable with this type of thinking. Individual atoms and molecules are difficult to observe, so often we must infer what they are doing from indirect evidence. Chemistry has matured tremendously during the past century by building an increasingly thorough understanding of atoms and molecules. We will introduce some of the basic concepts of atoms and molecules here and then refine these ideas as we progress through the text.

Chapter Objectives Online homework for this chapter may be assigned in OWL.

30

After mastering this chapter, you should be able to ❚ name at least three common polymers and give examples of their uses. ❚ define the terms atom, molecule, isotope, ion, compound, polymer, monomer, and functional group in your own words.

❚ describe the nuclear model for the atom and identify the numbers of protons, electrons, and neutrons in a particular isotope from its chemical symbol. ❚ calculate the atomic mass of an element from the masses and abundances of its isotopes. ❚ explain the difference between a molecular formula and an empirical formula. ❚ determine the number of atoms in a molecule from its chemical formula. ❚ describe the arrangement of elements in the periodic table and explain the usefulness of the table. ❚ obtain a correct chemical formula from a line drawing of an organic molecule. ❚ use standard chemical nomenclature to deduce the names of simple inorganic compounds from their formulas or vice versa. ❚ describe different forms of polyethylene and how their properties and applications are related to their molecular structures.

INSIGHT INTO

2.1

Polymers


As human civilization and technology have progressed, historical eras have been closely associated with materials from which important objects have been made. The Stone Age gave way to the Bronze Age, which in turn was followed by the Iron Age. These labels were chosen much later, through the lens of history, and it may be dangerous to try to characterize our own time period. But it isn’t hard to imagine that future archaeologists or historians might label the late 20th century and early 21st century as the Polymer Age. As you go about your life, the many plastics and synthetic fibers you encounter are examples of what chemists call polymers. The properties and applications of these polymers are so diverse that you may not even recognize that they have anything in common. Hard and durable plastics are routinely used as structural materials for things like computer cases and casual furniture. Softer, flexible plastics give us sandwich bags and Saran® wrap. Other polymers make up the nylon and rayon that are found in our carpets and clothing. Still more polymer materials, such as the filling in many bulletproof vests, offer incredible combinations of light weight and high strength. The diversity of polymer properties is truly impressive. And yet if we take the chemist’s approach and turn our attention to the atomic and molecular level, we will see that all of these polymers have much in common.

(a)

(b)

(c)

Figure 2.1 ❚ Polymers are the materials of choice for a host of everyday objects. All of the items shown in the left-hand photo are made from high-density polyethylene. The objects in the center are made from polystyrene, and those in the right-hand photo are made from poly(vinyl chloride).

2.1

Polymers

31

Because these molecules are so large, polymers are also called macromolecules.

Other uses of polyethylene are discussed in Section 2.7 at the end of this chapter.

Figure 2.2 ❚ Models showing how atoms are arranged in molecules of polyethylene, poly(vinyl chloride), and poly(vinylidene chloride).

Traced back to its Greek origins, the word polymer literally means “many parts.” That definition offers us a clue as to what these seemingly disparate substances all have in common. All polymers are made up of very large molecules. These large molecules are made up of many smaller molecules, linked end to end. A typical polymer molecule might contain hundreds or even thousands of these smaller constituent molecules, which are called monomers. If we look deeper into the composition of the polymer, we will find that these monomers are themselves made up of assemblies of atoms. But because polymer molecules are so large, it is often helpful to think of them as chains of monomers rather than collections of atoms. The observable macroscopic properties of any particular polymer depend on the identity of its constituent monomers, the number of monomers present, and the way that the monomers are connected to one another. To illustrate how dramatically the properties of a polymer can depend on its atomic composition, let’s consider three common items and look at the polymers from which they are made. The plastic bottles, in which your milk, juice, or shampoo come, are usually made of a plastic called polyethylene. Polyethylene molecules are composed entirely of just two elements, carbon and hydrogen. The carbon atoms are linked together in a long chain that is called the polymer backbone of the polymer molecule, and there are two hydrogen atoms attached to each carbon. The molecular model at the top of Figure 2.2 shows a portion of a polyethylene molecule. If one of the two hydrogen atoms on every other carbon is replaced with a chlorine atom, as shown in the middle of Figure 2.2, we will have poly(vinyl chloride). This polymer is commonly referred to as PVC. Plastic pipe made from PVC has been widely used in plumbing for many years, so you have

Molecules are depicted with balls representing atoms.

Hydrogen atoms are represented by white balls.

Carbon atoms are represented by black balls.

Polyethylene

Poly(vinyl chloride) Chlorine atoms are represented by green balls.

Poly(vinylidene chloride)

32

Chapter 2 Atoms and Molecules

probably seen PVC pipe in either your house or a hardware store. As you might guess, the PVC used for pipes is much harder and stronger than the polyethylene in soda bottles. Yet the chemical composition and structures of these two materials are very similar. Suppose that we replace the second hydrogen on every other carbon atom with chlorine, too. Then we will have poly(vinylidene chloride), which is the “plastic wrap” used in almost every kitchen to cover leftovers. These three common examples show how strongly the physical properties of a polymer are influenced by its chemical composition. To begin to explore the world of polymers systematically, we will need to know a little more about the atoms that are the simple building blocks of these giant molecules.

2.2

Atomic Structure and Mass

The comparison of polyethylene, poly(vinyl chloride), and poly(vinylidene chloride) demonstrates that the identity of the atoms in a molecule can have a tremendous impact on that molecule’s properties. Let’s begin our exploration by examining the structure of atoms, so that we can address the question of how a chlorine atom differs from a hydrogen atom. To do that, we’ll need to zoom in one more level, to the realm of subatomic particles.

Fundamental Concepts of the Atom Our current model of the structure of atoms has been accepted for nearly a century, but it took great creativity and many ingenious experiments to develop. The atom is composed of a small, compact core called the nucleus surrounded by a disperse cloud of electrons. The nucleus is composed of two types of particles: protons and neutrons. There is so much space between the electrons and the nucleus that it is impossible to show it to scale in an illustration. Consider Figure 2.3(a), which is similar to pictures you’ve seen before in high school chemistry or physical science books. The figure shows the relative positions of the protons, neutrons, and electrons. But if the protons and neutrons were actually the size shown, then the electrons would be hundreds of meters away. Another misunderstanding promoted by this type of illustration is the picture of electrons following regular orbits around the nucleus. A better model of atomic structure views the electrons as clouds of negative charge that surround the nucleus, as opposed to particles that orbit around it in an orderly way (Figure 2.3(b)). z

Orbit

Electron

y

x

Nucleus

(a)

We will examine the structure of atoms in much greater detail in Chapter 6.

(b)

Figure 2.3 ❚ Atoms have often been depicted as resembling a solar system: the nucleus is at the center, and the electrons orbit around it, as seen here in (a). Although such pictures do help to emphasize the way that protons, neutrons, and electrons are distributed in the atom, they cannot illustrate accurately the currently accepted model of atomic structure. Instead, we depict the electrons as clouds of negative charge surrounding the nucleus, as shown in (b). In such pictures, the density of the small dots is related to the probability of finding an electron at a particular location.

2.2 Atomic Structure and Mass

33

Protons and neutrons are themselves made up of even smaller particles, known as quarks.

We generally depict the charges in units of the electron charge, so that the charge of an electron is written as 1− and that of a proton is written as 1+.

Now we turn our attention to the numbers of protons, neutrons, and electrons in the atom. Electric charge provides an important constraint on these numbers. Protons are positively charged, electrons are negatively charged, and neutrons are neutral. Atoms themselves are also electrically neutral, so the numbers of protons and electrons present must be such that their charges will cancel each other. You may know from physics that the SI unit of charge is the coulomb (C). Experiments have shown that the electrical charges on a proton and an electron are equal and opposite. Every electron carries a charge of −1.602 × 10−19 C, whereas every proton carries a charge of +1.602 × 10−19 C. So for an atom to remain neutral, the numbers of electrons and protons must be equal. Because neutrons have no charge, the number of neutrons present is not restricted by the requirement for electrical neutrality. For most elements, the number of neutrons can vary from one atom to another, as we’ll see.

Atomic Number and Mass Number

The atomic mass unit is also referred to as a dalton and is sometimes abbreviated as u.

The number of protons in a particular atom, referred to as the atomic number, identifies the element. Carbon atoms make up the backbone of nearly all polymers, so we will consider them first. The atomic number of carbon is six, which tells us that a neutral carbon atom has six protons. Electrical neutrality requires that a carbon atom also must have six electrons. The great majority of carbon atoms—roughly 99%—also contain six neutrons. But some carbon atoms contain seven or even eight neutrons. Atoms of the same element that have different numbers of neutrons are called isotopes. Protons and electrons govern nearly all of the important chemical properties of atoms, so generally isotopes cannot be separated chemically. But the existence and even the relative abundance of isotopes can be proven by careful examinations of the mass of atoms. Protons and neutrons have similar masses; each is nearly 2000 times more massive than the electron. So the mass of any atom is concentrated in its nucleus. Individual atoms are so small and light that reporting their masses in conventional units such as kilograms or grams is not convenient. Instead we use a unit that is appropriate to the atomic scale: the atomic mass unit or amu. 1 amu = 1.6605 × 10−24 g Both the neutron and the proton have masses very close to one amu. The mass of a neutron is 1.009 amu, and that of a proton is 1.007 amu. The mass of an electron, in contrast, is just 0.00055 amu. So for many practical purposes, we can determine the mass of an atom simply by counting the number of protons and neutrons. That number will be the mass in amu, to a fairly reasonable approximation. Because of this, the combined total of protons and neutrons is called the mass number of the atom. Because isotopes are atoms of the same element with different numbers of neutrons, they will have the same atomic number but different mass numbers.

Isotopes How do we know that these isotopes exist? Modern instruments called mass spectrometers provide direct experimental evidence. The first important function of a mass spectrometer is to take a stream of microscopic particles—atoms or molecules—and “sort them” according to mass. (Figure 2.4 explains how the instrument does this.) Once the particles have been separated by mass, the second key function of the mass spectrometer is to measure accurately the number of particles with a given mass. The data are usually presented as a “mass spectrum.” Any time we refer to a spectrum, we will be noting a measurement that is made over a range of values of some variable. In this case that variable is mass, so the mass spectrum is really just a plot showing the number of particles detected as a function of mass. When a peak is seen at a particular mass, it means that the sample analyzed has some component with that mass. Figure 2.5 shows such a mass spectrum for a sample of carbon. Looking at the graph, we immediately see a large peak centered at mass 12. That represents the isotope 34


Magnet Beam of 12C+ ions

Accelerating plates

Collector Slit

–

Detector

+

Gas inlet

Beam of 4He+ ions

Electron gun

Figure 2.4 ❚ The schematic diagram shown here illustrates the key principles in the functioning of a mass spectrometer. A stream of gas to be analyzed enters at the left, and an electron gun causes some of the atoms to lose an electron, forming charged particles called ions. These ions are then accelerated to the right by an electric field, so that a beam of ions passes into a magnetic field. The magnetic field deflects the ions, and the extent of that deflection depends on the charge to mass ratio of each ion. For a given charge, lighter particles are deflected more severely than heavier ones. So if the sample contained both 4He+ and 12C+ ions, as shown here, the helium ions would be deflected much more than the carbon ions. This allows a slit to select ions of a particular charge to mass ratio, which then strike a detector. The current at this detector produces a signal that is proportional to the number of ions found with the desired charge to mass ratio, and this in turn is related to the amount of the parent gas molecule that entered the spectrometer.

Figure 2.5 ❚ A sketch of the mass

Signal

spectrum of elemental carbon is shown. The large peak is due to 12C, and the smaller peak to the right is 13 C. The size of the 13C peak here is somewhat exaggerated; it would actually be just 1/99 the size of the 12 C peak.

10

11

12 13 Mass (amu)

14

15

called carbon-12, whose nucleus contains six protons and six neutrons. This isotope is actually used to define the amu: an atom of carbon-12 has a mass of exactly 12 amu. But if we look at the mass spectrum closely, we also see a much smaller peak centered near mass 13. This tells us that there is a small amount of a second isotope, carbon-13, with seven neutrons. Comparing the relative sizes of the two peaks, we could determine that the carbon-12 isotope accounts for roughly 99% of the carbon atoms. More accurate measurement gives a value of 98.93%, with just 1.07% of carbon-13. It is also possible to determine that the exact mass of the carbon-13 isotope is 13.0036 amu. Because any sample we measure will always contain vast numbers of atoms, these same percentages, or isotopic abundances, will be found in any naturally occurring sample of carbon.

A number of radioactive isotopes of carbon are also known. The most common of these is 14C. Its abundance is measured in the carbon dating of archaeological objects.

Atomic Symbols All the information about the structure of the atom, which we have just discussed, can be written in scientific shorthand, using atomic symbols. The general atomic symbol can be written as A ZE


35

Table

❚

2.1

Names and symbols of some common elements whose symbols are not related to their English names

The term “ferrous metals” refers to iron or alloys such as steel that contain significant amounts of iron.

Name

Symbol (name origin)

Gold

Au (aurum)

Iron

Fe (ferrum)

Lead

Pb (plumbum)

Mercury

Hg (hydrargyrum)

Silver

Ag (argentum)

Sodium

Na (natrium)

Here E represents the atomic symbol for the element in question, the superscript A is the mass number, and the subscript Z is the atomic number. The symbol for carbon-12, for example, is 126C. Many atomic symbols are fairly obviously derived from the name of the element, such as the use of C for carbon in our example. For other elements, the symbol is based on the Latin name. The symbol for iron, for example, is Fe, derived from the Latin name ferrum. An atom of iron with 26 protons and 30 neutrons is represented as 56 26Fe. A listing of some common elements whose symbols are not based on their English names is provided in Table 2.1. A full list of elements and their symbols can be found in Appendix A at the back of this book.

Atomic Masses

6 C 12.011

Atomic number Symbol Relative atomic mass

Figure 2.6 ❚ Entry for carbon from a periodic table. The atomic number (6) and the atomic mass (12.011) are shown, along with the symbol for the element (C). Some tables may display additional information, and the exact layout may vary from one table to another. But once you are familiar with the table itself, usually it is easy to interpret whatever data are shown.

36

When you look at an entry in the periodic table, you see some of the information we’ve just defined, such as the atomic symbol and atomic number. Most periodic tables include additional information as well. Almost always, the atomic mass is given. This number provides the average mass in amu of an atom of the element. If you look up carbon in the periodic table inside the back cover of this book, you will find the box shown in Figure 2.6. The atomic mass appears under the symbol: 12.011. But we have already said that the mass of an atom of carbon-12 is exactly 12 amu, and that of carbon-13 is 13.0036 amu. So the value of 12.011 does not seem to be the mass of any individual atom of carbon. Then how are atomic masses defined and determined? The atomic mass is defined as the average mass of an atom of a particular element. Carbon has two stable isotopes with masses of 12.0000 and 13.0036 amu, respectively. So why is the average mass 12.011 and not something closer to 12.5? The answer is that when we take the average mass, we must account for the relative abundance of each isotope. Suppose that we could measure the mass of a 100-atom sample. Based on the isotopic abundances, we would expect to have 99 atoms of carbon-12 and only a single atom of carbon-13. In any sample that we can actually weigh, the number of atoms will be far greater than 100. Even using the best available laboratory balances, the smallest quantity of matter that can be weighed is about a nanogram, or 10−9 g. A nanogram of carbon would contain more than 1013 atoms. For such large numbers of atoms, it is safe to assume that the fraction of each isotope present will be determined by the natural abundances. For carbon, the fact that we only need to consider two stable isotopes makes the calculation fairly simple. We can


multiply the mass by the fractional abundance to weight each isotope’s contribution to the atomic mass. Carbon-12:

12.0000 × 0.9893 = 11.87

Carbon-13:

13.0036 × 0.0107 = 0.139

Weighted average mass = 11.87 + 0.139 = 12.01 The value of 12.011 found in the periodic table is obtained using additional significant figures on the isotopic abundance numbers.


2 .1

The chlorine present in PVC has two stable isotopes. 35Cl with a mass of 34.97 amu makes up 75.77% of the natural chlorine found. The other isotope is 37Cl, whose mass is 36.95 amu. What is the atomic mass of chlorine?

Strategy To determine the atomic mass, we must calculate the average mass weighted by the fractional abundance of each chlorine isotope. Because there are only two stable isotopes, their abundances must add up to 100%. So we can calculate the abundance of 37Cl from the given abundance of 35Cl.

Solution First, we calculate the abundance of the chlorine-37 isotope: Abundance of

Cl = 100% − 75.77% = 24.23%

37

Now we can calculate the contribution of each isotope to the atomic mass. 35

Cl:

34.97 × 0.7577 = 26.50

37

Cl:

36.95 × 0.2423 = 8.953

Weighted average mass = 26.50 + 8.953 = 35.45 So the atomic mass of chlorine is 35.45 amu.

Analyze Your Answer Based on the relative percentages, we should be able to decide if this answer makes sense. The individual isotopes have masses of roughly 35 and 37, so a 50/50 ratio would lead to an average mass of about 36. But the actual abundance of the 35Cl isotope is greater than that of 37Cl, so the average mass should be closer to 35. Thus our answer of 35.45 seems reasonable. And of course we can check the answer by consulting a periodic table. Discussion Some elements have several stable isotopes, but we can always do the same type of calculation by accounting for the mass and fractional abundance of each isotope.

Check Your Understanding There are three naturally occurring isotopes of the element silicon, which is widely used in producing computer chips. Given the masses and abundances below, calculate the atomic mass of silicon.

Isotope

Abundance

Mass

28

92.2%

27.977 amu

29

4.67%

28.977 amu

30

3.10%

29.974 amu

Si Si Si


37

2.3

The word cation is pronounced as “kat-ion,” not “kay-shun.”

Ions

In developing our ideas about the composition of atoms in the previous section, we used the fact that atoms are electrically neutral to conclude that the numbers of protons and electrons must be equal. When the number of protons and the number of electrons do not match, the result is a species with a net charge, called an ion. Such species are no longer atoms, and their behavior is markedly different from that of atoms. The operation of the mass spectrometer, illustrated in Figure 2.4, depends on the instrument’s ability to convert an atom into an ion. The separation of the particles by mass is generally based on the behavior of charged particles in a magnetic field, and the detector that counts the particles typically detects only ions, not neutral atoms. Ions can also play important roles in many chemical processes, including several that are important in the large-scale production of polymers. When an ion is derived from a single atom it is called a monatomic ion. When groups of atoms carry a charge they are called polyatomic ions. Monatomic or polyatomic ions may carry either negative or positive charges. Negatively charged ions are called anions, and they contain more electrons than protons. Similarly, an ion with more protons than electrons has a positive charge and is called a cation. We write symbols for ions analogously to those for atoms, adding the charge on the species as a superscript to the right of the atomic symbol. For monatomic ions, the number of protons still specifies the element whose symbol we use. Table 2.2 provides some examples of monatomic ions. Notice that the monatomic anions have names ending in -ide, whereas cations simply have “ion” added to the name of the element. The behavior and interaction of electrical charges are important topics in physics, but they also provide a basis for thinking about many aspects of chemistry. For our current interests, we will point out just two fundamental ideas about electric charge. First, opposite charges attract each other and like charges repel one another. And second, electric charge is conserved. These two ideas have important implications for the formation of ions in chemical processes. First of all, because charge is conserved, we can say that if a neutral atom or molecule is to be converted into an ion, some oppositely charged particle—most likely an electron or another ion—must be produced at the same time. Moreover, because opposite charges attract one another, some energy input is always needed to convert a neutral atom or molecule into a pair of oppositely charged particles.

Mathematical Description The statement that “opposites attract and likes repel” can be quantified mathematically. Coulomb’s law, which you may recall from a physics class, describes the

Table

❚

2.2

Examples of monatomic ions

Cation Name

38


Symbol

Anion Name

Symbol

Sodium ion

Na+

Fluoride ion

F–

Lithium ion

Li+

Chloride ion

Cl–

Potassium ion

K+

Bromide ion

Br–

Magnesium ion

Mg2+

Sulfide ion

S2–

Aluminum ion

Al3+

Nitride ion

N3–

Force

Figure 2.7 ❚ The figure shows how the coulombic force (Equation 2.1) varies with the distance r between two particles with opposite or like charges. When the charges have the same signs, the particles will repel one another, so the value of the force is positive. If the charges have opposite signs, the particles will attract one another and the value of the force will be negative.

Particles with the same charge have positive values, showing a repulsive force.

Distance between particles

0

Oppositely charged particles have negative values, showing an attractive force.

interaction of charged particles. The attraction of opposite charges and the repulsion of like charges are both described mathematically by one simple equation: q1q2 F = ————2 4pe0r

(2.1)

Here q1 and q2 are the charges, e0 is a constant called the permittivity of a vacuum, and r is the distance between the charges. F is the force the objects exert on one another as a result of their charges. Looking at this expression, when both charges have the same sign—either positive or negative—the resultant value for the force is a positive number. When the charges are opposite, the value is negative. This is consistent with the usual sign conventions used throughout chemistry and physics for force and energy; a negative value of F from Equation 2.1 indicates an attractive force and a positive value a repulsive one. Now consider the effect of varying the distance, r, between two ions. If two positively charged particles are initially very far apart (effectively infinite distance), the r 2 term in the denominator of Equation 2.1 will be very large. This in turn means that the force F will be very small, and so the particles will not interact with each other significantly. As the two like charges are brought closer together, the r 2 term in the denominator shrinks and so the (positive) force grows larger: the particles repel each other. If we somehow force the particles closer together, the repulsive force will continue to grow. The distance dependence of the coulombic force is illustrated in Figure 2.7.

Ions and Their Properties Many monatomic cations and anions exist. These ions can exist in the gas phase, and many are important in atmospheric chemistry. But we encounter ions most frequently when dealing with the chemistry of substances dissolved in water. For example, sodium atoms lose an electron relatively easily to form the sodium cation, Na+. Because it still has 11 protons, this ion retains the symbol of sodium, yet it does not behave at all like an atom of sodium. Consider an order of French fries. You may have heard news stories about the high amount of sodium in an order of fries, and concerns have been raised about the possible health effects of too much sodium in our diets. This statement could be confusing because here the word “sodium” does not refer to sodium metal. In fact, if we place sodium metal on freshly made French fries, the metal will burst into flame! The sodium we hear about in stories on diet and health is actually sodium ion, which is added to the fries when salt is sprinkled on. Too much salt might still be a health concern, but we certainly don’t worry about the salt igniting. There is a big difference between ions and atoms, at least in this case. In contrast to sodium, chlorine readily gains an extra electron forming the chloride ion Cl−. Again, there is a noticeable difference between the ion and the atom of

2.3

Ions

39

Chlorine gas released in a January 2005 train accident in South Carolina led to eight deaths and forced many residents from their homes for days.

chlorine. The table salt we discussed above is sodium chloride, which contains chloride anions. Just like sodium, these chloride ions are present in French fries or any other salted foods. Chlorine atoms, on the other hand, combine in pairs to form a yellowish-green gas, Cl2, which irritates the lungs and can be toxic. The behavior of the ion is clearly much different from that of the neutral atom or molecule. Polymers do not generally contain ions. But ions are important in the chemical reactions used to produce many common plastics. As a result of their electric charge, ions are often much more reactive than neutral atoms or molecules. So small amounts of ions are often used to initiate or sustain the chemical reactions that join monomers together to form polymers.

2.4

Compounds and Chemical Bonds

A basic picture of atoms is a good starting point for understanding the properties of polymers. But to begin to see how the observable properties of a polymer might be related to its atomic and molecular makeup, we will need to consider the connections between atoms. Which atoms are actually attached to one another? And what different types of connections—or chemical bonds—are involved? Once again we’ll begin by trying to establish some vocabulary that will help us understand compounds and the chemical bonds that hold them together.

Chemical Formulas

In Section 3.5, we will learn how to find the empirical formula for a compound from experimental data.

Ethylene, C2H4

40

A chemical compound is a pure substance made up of atoms of two or more elements joined together by chemical bonds. In any compound, the atoms combine in fixed whole number ratios. In any such combination, the resultant substance behaves differently from the atoms alone. In many compounds, atoms combine to form discrete particles called molecules. Molecules can be broken down into their constituent atoms, but the resulting collection of atoms no longer behaves like the original molecule. Other materials are composed of vast arrays or extended structures of atoms or ions but do not form discrete molecules. Alloys, metals, and ionic solids (composed of paired ions) fall into this category of chemical compounds. We’ve seen how we can use atomic symbols as shorthand notation to designate atoms. That same idea can be extended to describe the composition of either molecules or extended compounds in a simple symbolic representation. A chemical formula describes a compound in terms of its constituent elements. We will actually encounter two distinct types of chemical formulas: molecular formulas and empirical formulas. The molecular formula of a compound is a kind of parts list that describes the atomic composition of a molecule efficiently. The molecular formula of the ethylene monomer from which polyethylene is produced is C2H4; this tells us that there are two carbon atoms and four hydrogen atoms per molecule. The empirical formula tells us only the relative ratio between the numbers of atoms of the different elements present. Let’s consider ethylene again. The ratio of carbon atoms to hydrogen is 1:2. So the empirical formula is CH2. When dealing with an empirical formula, it is important to realize that it does not tell how large or small an individual molecule of the compound might be; only the relative numbers of atoms of each element are given. We often emphasize this fact by writing a subscript ‘n’ on the entire formula. For ethylene, this would give us (CH2)n, which means that each molecule must contain some integral number of CH2 units. Empirical formulas, or a minor variation on them, are especially common when dealing with polymer molecules. Because polymer molecules are so large, the exact number of monomer units in a molecule is generally not very important. And in fact, the exact length of the polymer chains is often not the same for all molecules in a given sample. Instead, there is usually some range of chain lengths that will exist,


depending on how the polymer was actually produced. As long as the chains are all within some reasonable range of lengths, the macroscopic properties of the polymer are not affected substantially. So polymer formulas are most often written like empirical formulas. The repeating unit contributed by each monomer molecule is written in parentheses or brackets, and a subscript ‘n’ is used to emphasize that a large number of these units will be found in any individual molecule. For polyethylene, we would write the formula as ![CH2CH2]n!. Here the dashes are added to stress that these units are attached end to end to build up the long chain of the polymer. For the most common forms of polyethylene, the number of monomer units (i.e., the value of ‘n’) is in the tens of thousands. We could write similar formulas for the other polymers mentioned in the opening section on pages 32 and 33. Poly(vinyl chloride): Poly(vinylidene chloride):

In the closing Insight section for this chapter, we will look more closely at the way large changes in chain length influence the properties of polyethylene.

![CH2CHCl]n! ![CH2CCl2]n!

There are four rules that allow us to write most formulas that we will need in this textbook. 1. Indicate the types of atoms in the substance by their atomic symbols. 2. The number of each atom in the compound is indicated by a subscript to the right of the atomic symbol. For example, the chemical formula of ethylene, C2H4, tells us that each molecule contains two carbon atoms and four hydrogen atoms. 3. Groups of atoms can be designated by using parentheses. Subscripts outside these parentheses mean that all atoms enclosed in the parentheses are multiplied by the value indicated in the subscript. 4. Water molecules associated with certain compounds called hydrates are indicated separately from the rest of the compound. Example Problem 2.2 shows how to interpret chemical formulas by inverting some of these rules.


2.2

We cannot generally produce a polymer by simply mixing a large sample of the desired monomers. Instead, additional compounds called initiators or catalysts are almost always needed to start a polymerization. One polymerization catalyst is diethylaluminum chloride, Al(C2H5)2Cl. How many of each type of atom are in a molecule of this compound?

Strategy The subscripts in a formula indicate how many atoms of each type are in the molecule. The parentheses designate a group of atoms, and the subscript associated with the parentheses multiplies each atom in the group.

Solution In each molecule of Al(C 2H 5) 2Cl, there is one aluminum atom, one chlorine atom, and two groups of C 2H 5. Each of the C 2H 5 groups contains two carbon atoms and five hydrogen atoms. We multiply those numbers by two because there are two C2H5 groups present; so we have four carbon atoms and ten hydrogen atoms.

Discussion The number of atoms present might be easier to see if we wrote this formula as AlC4H10Cl. Right now you might feel that would be simpler. But when we write it as Al(C2H5)2Cl, we are actually conveying some additional information about the way the atoms are connected to one another. Specifically, we are showing that the carbon and hydrogen atoms are arranged as two C2H5 groups and that each of these groups is attached to the aluminum atom. Later we will learn that such a C2H5 unit is called an ethyl group.


41

Check Your Understanding A compound with the rather imposing name of 2,29-azo-bis-isobutyrylnitrile is used to initiate the growth of some polymers, including poly(vinyl chloride). If the molecular formula is C8H12N4, how many of each type of atom are in a molecule of the compound? What is the empirical formula of this compound?

Chemical Bonding

Lattices are often depicted as having shapes that are essentially cubic, but there are actually 17 different geometric shapes, not all of which are cubic.

Figure 2.8 ❚ A conceptual illustration showing the transfer of one electron from a sodium atom to a chlorine atom, forming a pair of ions (Na+ and Cl–). Once electron transfer takes place, coulombic force draws the ions together.

Atoms combine to make compounds by forming chemical bonds. Several different types of chemical bonds are possible, and once we learn to recognize them, these types of bonds will help us to understand some of the chemical properties of many substances. All chemical bonds share two characteristics. First, all bonds involve exchange or sharing of electrons. We will return to this concept often in this text as we investigate chemical reactions and properties of molecules. Second, this exchange or sharing of electrons results in lower energy for the compound relative to the separate atoms. A chemical bond will not form, or will have only a fleeting existence, unless it lowers the overall energy of the collection of atoms involved. Chemical bonds can be divided into three broad categories: ionic, covalent, and metallic. Some compounds are composed of collections of oppositely charged ions that form an extended array called a lattice. The bonding in these compounds is called ionic bonding. To form the ions that make up the compound, one substance loses an electron to become a cation, while another gains an electron to become an anion. We can view this as the transfer of an electron from one species to another. Figure 2.8 shows this concept for one ionic compound, NaCl. Ionic compounds form extended systems or lattices of alternating positive and negative charges, such as that shown in Figure 2.9. Although the formula NaCl correctly indicates that sodium and chlorine are present in a 1:1 ratio, we cannot really identify an individual “molecule” of NaCl. To emphasize this distinction, we sometimes refer to a formula unit, rather than a molecule, when talking about ionic

The dots represent outer electrons.

Na

Cl Transfer of an electron produces a pair of oppositely charged ions.

Cl

⫺

⫹ Na

Cl

Oppositely charged ions attract one another.

⫺

⫹ Na

42


Cl

Cl–

Na+

Cl–

Na+

compounds. The formula unit is the smallest whole number ratio of atoms in an ionic compound. Metals represent another type of extended system, but here the chemical bonding is totally different. In metals, the atoms are once again arranged in a lattice, but positively and negatively charged species do not alternate. Instead, the nuclei and some fraction of their electrons comprise a positively charged “core” localized at these lattice points, and other electrons move more or less freely throughout the whole array. This is called metallic bonding. Metallic bonding leads to electrical conductivity because electrons can move easily through the bulk material. Figure 2.10 shows a schematic illustration of the concept of metallic bonding. When electrons are shared between pairs of atoms rather than donated from one atom to another or mobile across an entire lattice, we have covalent bonds. In covalent bonds, electrons are usually shared in pairs. Two electrons (and sometimes four or six) are located between two nuclei and the sharing leads to an attraction between the nuclei. The long chains in all polymers are formed by covalent bonds in which electrons are shared between adjacent carbon atoms. Smaller, more familiar molecules such as water, carbon dioxide, and propane are simpler examples. All three types of chemical bonds will be discussed in much greater detail in Chapters 7 and 8. Polymer molecules are built up by the successive addition of monomers to form characteristic long-chain backbones. The bonds that hold the monomers to one another, as well as the bonds between atoms within each monomer unit, are covalent bonds. But ionic bonding is important in many compounds that are used to initiate or sustain the reactions needed to grow a polymer.

Figure 2.9 ❚ Two different representations of the NaCl crystal structure are shown. In each case, the green spheres represent chloride anions, and the gray spheres denote sodium cations. The view on the left emphasizes the positions of the ions, and that on the right better illustrates their relative sizes. In a macroscopic salt crystal, additional ions would simply extend this structure, repeating the same alternating pattern.

Water, H2O

Carbon dioxide, CO2

Propane, C3H8

The nucleus and inner electrons provide a positively charged “core.”

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

The outer electrons form a “sea” of negative charge surrounding the positive cores.

Figure 2.10 ❚ In this simple conceptual picture of metallic bonding, each metal atom contributes one or more electrons to a mobile “electron sea.” The ability of the electrons to move freely through this “sea” allows the metal to conduct electricity. Here the blue area depicts those mobile (or “delocalized”) electrons, and the red circles represent the positively charged “cores” of the individual atoms.


43

2.5

One of the difficulties that persisted in the development of the periodic table was the controversy over assigning atomic masses. Without a unique scale for atomic mass, scientists drew different conclusions about how to arrange the elements “in order.”

The Periodic Table

One of the most recognizable tools of chemistry is the periodic table. It is prominently displayed in practically every chemistry classroom, and many chemists own T-shirts, neckties, or coffee mugs on which the periodic table appears as a professional badge of honor. Why do chemists hold this table in such high regard? Once you are familiar with it, the periodic table summarizes a wealth of information about the behavior of elements, organizing them simultaneously in ascending order of atomic number and in groups according to chemical behavior. An experienced chemist can get a rough idea of an element’s properties simply from where that element sits in the periodic table. Today we can purchase or download artistic renderings of the periodic table, color-coded to display properties of particular interest for different applications. But like many developments in science, the emergence of what we now accept as the periodic table was accompanied by some degree of controversy. A number of scientists had devised various schemes for arranging the elements. These attempts to organize the understanding of the elements were not well received, however. One proposal by John Newlands in 1866 would have grouped elements similarly to musical octaves. This idea was literally laughed at during a scientific meeting where one critic sarcastically asked whether a listing by alphabetical order had been tried, noting, “any arrangement would present occasional coincidences.” Despite the skepticism of the 19th century scientific community, efforts to organize the elements persisted. Numerous observations suggested a regularity, or periodicity, in the behavior of the elements known at that time. By 1869, Russian scientist Dmitri Mendeleev had published his first periodic table and enumerated the periodic law: when properly arranged, the elements display a regular and periodic variation in their chemical properties. The most significant and impressive feature of Mendeleev’s work was his prediction of the existence of undiscovered elements. He left holes in his proposed table at positions where no known element seemed to fit. Later, when the elements to fill in these holes were identified, the scientific community accepted Mendeleev’s work. The discovery of the periodic law and construction of the periodic table represents one of the most significant creative insights in the history of chemistry. Prior to Mendeleev’s time, chemists had to learn the properties of each element individually. As more and more elements were discovered, that task became increasingly daunting. The periodic table helped the study of chemistry to expand quickly by providing a simple, visual means to organize the elements in terms of their chemical and physical properties.

Periods and Groups The modern periodic table simultaneously arranges elements in two important ways: the horizontal rows of the table, called periods, and the vertical columns, called groups. The term “period” is used for the rows because many important properties of the elements vary systematically as we move across a row. Figure 2.11 shows a plot of the density of elements, all in their solid state, as a function of atomic number. From the graph, it is clear that density varies according to a fairly regular pattern that goes through a series of minima and maxima. Different colors are used for the data points in this graph to show how the variation in density is correlated with position in the periodic table. Each color represents a period (row) in the table. Because the elements in the periodic table are arranged in order of increasing atomic number, moving across each segment of the graph corresponds to moving from left to right across the corresponding row of the periodic table. You can see readily that as we move across a row in this way, the density of the elements is initially small, increases until passing through a maximum, and then decreases again. Figure 2.12 shows the same data, with the density represented by the shading of each element’s box. This representation clearly shows how the density of the elements varies regularly across

44


Figure 2.11 ❚ The density of

25000

Density (kg/cubic meter)

20000

15000

10000

5000

0 0

10

20

30

40 50 60 70 Atomic number

80

90 100

each row of the table. The rows in the table are numbered 1 through 7 sequentially from top to bottom. Although the properties of the elements can vary widely across a period, each column collects elements that have similar chemical properties. Most elements can combine with hydrogen to form compounds. The graph in Figure 2.13 shows the number of hydrogen atoms with which an atom of each element will combine, and the regular variation in the plot clearly shows that this is a periodic property. Elements in a group (column) combine with the same number of hydrogen atoms. Fluorine, chlorine, and bromine each combine with one atom of hydrogen, for example, and all fall in the same group. These types of chemical similarities were among the evidence that led to the development of the periodic table, so some of the groups predate the general acceptance of the table. These groups of elements were assigned names and those names have remained with them. Thus the elements in the far left-hand column (Li, Na, K, Rb, and Cs) are known collectively as alkali metals. Similarly, Be, Mg, Ca, Sr, and Ba are called alkaline earths, and F, Cl, Br, and I are referred to as halogens. He, Ne, Ar, Kr, and Xe were discovered much later than most of the other elements, and they have been named rare gases or noble gases. Other groups are named, but their names are less commonly used and won’t be mentioned here.

1 H Li

17

18

H

He

2

13

14

15

16

Be

B

C

N

O

F

Ne

Al

Si

P

S

Cl

Ar

Na Mg

3

4

5

6

7

V

Cr

8

9

10

11

12

Mn Fe

Co

Ni

Cu

Zn

Ga Ge

As

Se

Br

Kr

K

Ca

Sc

Ti

Rb

Sr

Y

Zr

Nb Mo

Tc

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

I

Xe

Cs

Ba

La

Hf

Ta

W

Re

Os

Ir

Pt

Au Hg

Tl

Pb

Bi

Po

At

Rn

Fr

Ra

Ac

Rf

Ha

Sg

Ns

Hs

Mt

Ce

Pr

Nd Pm Sm

Eu Gd

Tb Dy

Ho

Er

Tm Yb

Lu

Th

Pa

U

Am Cm

Bk

Es

Fm Md No

Lr

Np Pu

Cf

elements in their solid states is plotted as a function of atomic number. Here density is in units of kg/m3. The various colors represent the different periods (rows) in the periodic table. Notice how the same general pattern repeats as we move across each row: the density is low at the left-hand edge of the row (Group 1), increases through a maximum, and then decreases as we continue to move from left to right across the table. (Only those elements that exist as solids under ordinary conditions are shown.)

The noble gas elements were once called inert gases because they were considered completely unreactive. Several noble gas compounds are now known, so the term inert is no longer used.

Figure 2.12 ❚ The data from Figure 2.11 are presented here in a different form. The shading of the boxes in the periodic table represents the density of each element; darker shading indicates higher density. The general trends in density are apparent as you look across a row or down a column.

2.5 The Periodic Table

45

Number of hydrogens bonded to element

Figure 2.13 ❚ The graph shows the number of hydrogen atoms with which an individual atom of various elements will combine. The periodicity of this chemical property is evident from the cyclic graph.

4

3

2

1

0 1

3

5

7

9 11 13 Atomic number

15

17

There are also names for different regions of the table. Elements in the two groups on the left side of the table and the six groups on the right side are collectively referred to as representative elements, or main group elements. Elements that separate these two parts of the representative groups in the main body of the periodic table are called transition metals. Iron is an example of a transition metal. The elements that appear below the rest of the periodic table are called lanthanides (named after the element lanthanum, Z = 57) and actinides (named after the element actinium, Z = 89). In addition to these names, several numbering systems have been used to designate groups. Current convention dictates numbering from left to right starting with 1 and proceeding to 18. Thus, for example, the group containing C, Si, Ge, Sn, and Pb is referred to as Group 14.

Metals, Nonmetals, and Metalloids As we move across the periodic table, atomic properties vary gradually rather than abruptly. So the boundaries between these categories are ambiguous in some ways.

46

Another way to classify an element is as a metal, nonmetal, or metalloid. Once again, the periodic table conveniently arranges elements so that one can place a given element easily into one of these categories. Most of the elements are metals. Their general location in the periodic table is toward the left and bottom, as seen in the coloring of the periodic table in Figure 2.14. Metals share a number of similarities in chemical and physical properties. Physically, metals are shiny, malleable, and ductile (meaning they can be pulled into wires). They also conduct electricity, so wires are always made from metals. Chemical properties can also be used to distinguish metals. Metallic elements tend to form cations in most of their compounds, for example. Nonmetals occupy the upper right-hand portion of the periodic table. There are fewer nonmetals than metals. But when we consider the relative importance of elements, nonmetals hold their own because of their role in the chemistry of living things. Most of the molecules that make up the human body consist predominantly or exclusively of the nonmetallic elements carbon, hydrogen, oxygen, nitrogen, sulfur, and phosphorus. As our examples so far might lead you to guess, polymers also consist almost exclusively of nonmetallic elements. In contrast to metals, nonmetals are not shiny, malleable, or ductile nor are they good conductors of electricity. These physical properties provide one means by which we can distinguish metals from nonmetals. Whether an element is a metal or nonmetal may seem simple to determine based on the physical properties cited above. Yet, some elements cannot be classified easily as either metal or nonmetal. The question whether or not a substance conducts electricity, for example, does not always have a simple yes or no answer. Lacking a reliable means of drawing a clean boundary between the two categories, scientists have generally chosen to refer to intermediate cases as metalloids or semimetals. In the periodic table, metalloids are clustered along a diagonal path, as shown in Figure 2.14. This categorization


1

17

1

1

2

H

He

H

2

3

4

Li

Be

11

12

Metals Metalloids Nonmetals

Na Mg

13

14

15

5

6

7

8

9

10

B

C

N

O

F

Ne

13

14

15

16

17

18

Al

Si

P

S

Cl

Ar

3

4

5

6

7

8

9

10

11

19

20

21

22

23

24

25

26

27

28

29

30

31

32

33

34

35

36

K

Ca

Sc

Ti

V

Cr

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

Kr

37

38

39

40

41

42

Rb

Sr

Y

Zr

Mn Fe 43

Nb Mo Tc

12

16

18

44

45

46

47

48

49

50

51

52

53

54

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

I

Xe

55

56

57

72

73

74

75

76

77

78

79

80

81

82

83

84

85

86

Cs

Ba

La

Hf

Ta

W

Re

Os

Ir

Pt

Au

Hg

Tl

Pb

Bi

Po

At

Rn

88

89

87

Fr

Ra Ac

104

105

106

107

108

109

Rf

Ha

Sg

Ns

Hs

Mt

58

59

Ce

Pr

60

61

62

Nd Pm Sm

90

91

92

93

Th

Pa

U

Np

94

63

64

65

Eu

Gd

Tb

95

96

97

Pu Am Cm Bk

66

67

Dy Ho 98

99

Cf

Es

68

Er 100

69

70

Tm Yb 101

71

Lu

102

103

Fm Md No

Lr

Figure 2.14 ❚ The colors in this periodic table identify each element as a metal, nonmetal, or metalloid. Notice how the metals are grouped toward the left and bottom of the table, and nonmetals are toward the upper right.

gives us some useful flexibility and also emphasizes that properties change gradually rather than abruptly as one moves across or down the periodic table. All of the polymer molecules we have mentioned are carbon based. Their skeletons consist entirely of carbon atoms. Because elements in the same group of the periodic table have similar chemical behavior, you might be wondering whether similar polymers could be produced based on silicon, which appears right below carbon in Group 14. Silicon-based polymers, known as silicones, do exist, but they differ from carbon polymers in important ways. Covalent bonds can be formed between silicon atoms, but they are not as strong as those between carbon atoms. So chains of silicon atoms become unstable beyond a length of around ten atoms, and silicon analogs of polymers such as polyethylene cannot be produced. Instead of pure silicon, the backbone chains in silicone polymers consist of alternating silicon and oxygen atoms. The Si!O bond is strong enough to allow these chains to grow quite long. Additional atoms or groups of atoms bound to the silicon atoms influence the properties of the polymer. The range of polymers that can be produced is not nearly as diverse as for carbon, but silicone polymers are widely used in applications including greases, caulking materials, water repellents, and surfactants.

2.6

Proteins are natural polymers whose backbones include carbon, oxygen, and nitrogen. The specific arrangement that allows this type of backbone is called a peptide bond.


Just as engineering can be broken down into various specialties, so too, chemistry can be viewed as a collection of subfields. Two of the most fundamental areas of chemistry are organic chemistry and inorganic chemistry. These names arise from the fact that at one time organic chemistry would have been defined as the chemistry of living things. A more modern definition is that organic chemistry is the study of compounds of the element carbon. As we’ve already seen, this includes naturally occurring biological molecules and also nearly all synthetic polymers. Inorganic chemistry is the study of all other elements and their compounds. At first glance, this may strike you as a

The division of chemistry into subdisciplines is a historic artifact. Much modern research in chemistry takes place at the interface of two or more subdisciplines.


47

bizarre distinction because carbon is just one element in a periodic table that contains more than 100 others. But the chemistry of carbon is so rich, diverse, and important that organic chemists actually outnumber inorganic chemists. Because this text is intended as a brief overview of important chemical principles, we will focus most of our attention on the behavior of molecules in general rather than delving too deeply into the details of either subfield. In this section, we will describe briefly some important similarities and differences between organic and inorganic chemistry and introduce terminology and notation that we will need as we progress through the text.

Inorganic Chemistry—Main Groups and Transition Metals We will study chemical bonding and molecular shapes in detail in Chapter 7.

The solid and dashed triangles in the structure on the left indicate that one of the chlorine atoms would be in front of the plane of the page and one would be behind that plane.

Many inorganic compounds exist as relatively small molecules whose atoms are joined together through covalent bonds. One such compound is silicon tetrachloride, SiCl4, which has important uses in the production of semiconductors. Figure 2.15 shows some visual representations of SiCl4 that we might use to illustrate how its atoms are actually arranged into molecules. Four chlorine atoms surround a central silicon atom, and each chlorine shares one pair of electrons with the silicon. Silicon and chlorine are both main group elements, found in Groups 14 and 17 of the periodic table, respectively. As mentioned in the previous section, elements from the same group tend to display similar chemical properties. Thus once we know that SiCl4 exists, we might expect that other pairs of elements from the same groups might form similar compounds. And this prediction is correct: compounds such as SnCl4 and CF4 do exist and have structures and bonds analogous to those in Figure 2.15 for SiCl4. Other compounds of the main group elements form extended ionic structures, such as that of NaCl in Figure 2.9. But despite the difference in the types of chemical bonds employed, we can still readily predict that similar compounds should exist for other pairs of elements from the same groups. From the periodic table, we see that sodium is in Group 1 and chlorine is in Group 17. So we can expect that other pairs of elements from these columns of the table will form ionic solids, too. Again, our prediction is accurate; compounds such as LiCl, NaF, and KBr have structures analogous to that of NaCl. The reason for the existence of these similar compounds is simple. All of the metals in Group 1 form cations with a 1+ charge, and all of the elements in Group 17 form anions with a 1− charge. Any of these cations can combine with any of the anions in a 1:1 ratio to form neutral compounds. The chemistry of transition metals is somewhat more complicated than that of the main group elements, though, because most transition metals can form multiple cations with different charges. Iron commonly forms two different monatomic cations: Fe2+ and Fe3+. As a result of this, iron can form a more diverse set of compounds than Group 1 metals. It can combine with chlorine to form either FeCl2 or FeCl3, and these two compounds have significantly different physical properties (Figure 2.16). Largely because they can form multiple cations, the chemistry of transition metals does not

Cl

Cl Si Cl

Cl

Figure 2.15 ❚ This figure presents three depictions of SiCl4. In the drawing at the left, each atom is represented by its symbol and the lines between the symbols depict chemical bonds. In the center panel is a “ball and stick” model, where each atom is a ball, and the bonds are shown as sticks connecting the balls. In the right-hand panel is a “space filling” model, where atoms are shown as balls that overlap one another strongly. Each type of model is commonly used, and each has its strengths and weaknesses. 48



vary as sharply from group to group. Regardless of their positions in the periodic table, for example, most transition metals can form cations with a 2+ charge. Thus predictions based simply on group number are not as reliable here as they are for the representative elements. When considering transition metals and their compounds, we must rely more heavily on knowledge of the specific chemistry of each element.

Organic Chemistry


Iron(III) chloride, FeCl3 (Here forming as the solid at the bottom of the test tube) Orange-brown color Density 2.90 g cm-3 Melts at 306°C

© Cengage Learning/ Charles D. Winters

All organic compounds feature carbon skeletons, similar to those we have already seen in our introduction to polymers. Other elements frequently found in organic compounds include hydrogen, oxygen, and nitrogen. Despite this rather short list of elements, more than 18 million organic compounds exist. This vast number of compounds arises from some unusual aspects of the chemistry of carbon itself. Most importantly, carbon atoms readily attach to one another to form chains, and these chains can grow quite long. Many of the polymer molecules we have been discussing in this chapter contain thousands of carbon atoms. Furthermore, some of these long chains are straight, whereas others are branched at one or more places. And finally, any pair of carbon atoms can bond to one another in three different ways, by sharing either one, two, or three pairs of electrons. When taken all together, these factors allow carbon to form a vast array of compounds. The diversity of organic compounds presents some challenges. It is not uncommon for several different organic compounds to have the same molecular formula, for example, but to display different properties depending on exactly how the atoms are arranged into molecules. (Different compounds with the same molecular formula are called isomers.) So organic chemists frequently must depict molecules not only by their formulas, but in some way that also conveys important information about the arrangement of the atoms. This could be done using structural formulas of the sort we saw in Figure 2.15. But because organic chemistry often deals with very large compounds and complex structures, that option is somewhat unwieldy. A shorthand notation known as a line structure has emerged as the most common method for describing organic compounds simply and unambiguously. The line structure is a modified version of the structural formula. As in any structural formula, lines are used to depict bonds between atoms. But in a line drawing, many of the elemental symbols are omitted. By definition, an organic compound is based on carbon atoms. So to reduce clutter in a line drawing, the ‘C’ symbols for carbon atoms are not written. Furthermore, because organic compounds almost always contain many hydrogen atoms, the ‘H’ symbol for any hydrogen atom that is attached directly to a carbon atom is also not written. Symbols are written for any elements other than carbon and hydrogen, as well as for any hydrogen atoms that are not directly attached to carbon. We can illustrate the relationship between a structural formula and a line drawing with an example.

Iron(II) chloride, FeCl2 Greenish-yellow color Density 3.16 g cm-3 Melts at 670°C

Figure 2.16 ❚ Transition metals typically form more than one type of cation, giving them a very diverse chemistry. Iron, for example, forms cations with both 2+ and 3+ charges, and this allows it to form two different ionic compounds with chlorine. FeCl2 and FeCl3 have different appearances and properties.

2.3

Poly(methyl methacrylate) is widely known as Plexiglas®. The structural formula for the monomer, methyl methacrylate, is shown below. Write the corresponding line structure for methyl methacrylate. H

H C H

H C H

C

C O

O

C

H

HH

Strategy We convert the structure into a line drawing by removing the symbols for all carbon atoms and for hydrogen atoms attached directly to carbons. Bonds to or between 2.6 Inorganic and Organic Chemistry

49

carbon atoms remain, so carbon atom positions become either intersections between lines or ends of lines. Bonds between carbon and hydrogen atoms are omitted.

Solution First, we will remove the symbols and bonds for all of the hydrogen atoms because they are all bound directly to carbon. C C

C

C

O

C

O Next we remove the symbols for the carbon atoms, leaving intact the lines that depict the remaining bonds. This gives us the final line structure.

O O

Discussion The line structure is much more compact than the original structural formula. An experienced chemist quickly recognizes where the atoms whose symbols are not shown need to be.

Check Your Understanding The structural formula for styrene, which is the monomer for the common plastic polystyrene, is shown below. Convert this to a line drawing. H H H

H

C C

C C

C

C

C C

H H

H

H

We will use these line structures throughout the rest of this textbook, and you may also encounter them in other places, such as the information sheets that accompany prescription drugs. In many instances, it will be necessary to interpret the line drawing to determine the molecular formula, so we should develop a way to do that systematically. In addition to the rules we used before to transform a structural diagram into a line structure, we will need to introduce two important generalizations about chemical bonding. 1. A hydrogen atom in an organic molecule can form only one covalent bond to one other atom. 2. Every carbon atom in an organic molecule always forms exactly four covalent bonds. In combination, these two facts allow us to fill in all of the carbon and hydrogen atoms that are not explicitly written in a line structure. First, we place a carbon anywhere that we see either an intersection between lines or the end of a line. Then we add hydrogen atoms as needed to bring each carbon’s number of bonds up to four.

50



2.4

A temperature resistant plastic called poly(phenylene oxide) is a key component of resins such as GE’s Noryl®, which is widely used in computer cases and automobile dashboards. The line structure below represents 2,6-dimethylphenol, which is the monomer from which poly(phenylene oxide) is made. What is the molecular formula for 2,6-dimethylphenol? OH

Strategy First, we can pencil in carbon atoms at the appropriate positions. Then we will add hydrogen atoms as needed. Once all of the atoms have been identified, it will be easy to count them to produce the needed formula.

Solution We place a carbon atom at the end of a line or the intersection between lines. OH C

C C

C

C

C

C

C

Next we count the number of bonds shown for each carbon. If that number is three or less, we add as many hydrogen atoms as needed to bring it up to four. OH H3C

C HC

C

C H

C

CH3

CH

Now all of the atoms are shown explicitly. Counting, we arrive at the molecular formula as C8H10O.

Discussion The double lines in the ring in this structure represent double bonds, in which two pairs of electrons are shared between two atoms. Notice that in locating the carbon atoms, we treated the double lines the same as we did the single lines: each intersection represents a carbon atom, no matter how many lines meet.

Check Your Understanding Once, poly(vinylpyrrolidone) was used in the manufacture of hairsprays, and it is still used in the glue that holds the layers together in plywood. The line structure for the vinylpyrrolidone monomer is shown below. Find the corresponding molecular formula.

N

O


51

Functional Groups

The role of addition reactions in producing polymers will be examined in Section 2.8.

Not all organic compounds that contain !OH groups are alcohols. Carboxylic acids contain a !COOH functional group, for example.

Given the vast number of organic compounds, the need for some systematic way to understand their chemistry should be apparent. One of the most important concepts for an organic chemist is the idea that certain arrangements of atoms tend to display similar chemical properties whenever they appear together. Such an arrangement of atoms is called a functional group. One of the simplest functional groups, and one that is central to many polymerization reactions, is a pair of carbon atoms joined by a double bond. If the double bond is converted to a single bond, then each carbon atom can form a new bond to another atom. Thus, the characteristic reaction of a carbon-carbon double bond is addition, in which new atoms or groups of atoms are attached to a molecule. Line structures make it very easy to identify any C"C groups in a molecule, and thus to locate positions at which addition reactions might be feasible. The simplest organic compounds are hydrocarbons, molecules that contain only carbon and hydrogen atoms. We can imagine the formation of more complicated molecules by replacing one or more of a hydrocarbon’s hydrogen atoms with a functional group. Compounds in which a hydrogen atom is replaced by an !OH functional group, for example, are collectively referred to as alcohols. The presence of the !OH group conveys certain properties to this class of molecules, including the ability to mix with water to a much greater extent than the corresponding hydrocarbons. Often, the notion of functional groups influences the way in which we choose to write chemical formulas. If the chemical formula for an alcohol is written so that the !OH group is emphasized, then it will be easier to recognize that this group is present. So the formula for the simplest alcohol, methanol, is most often written as CH3OH rather than as CH4O. Similarly, ethanol is generally written as C2H5OH rather than C2H6O. Other common functional groups are listed in Table 2.3.

Table

❚

2.3

Some common functional groups

Functional Group

Class of Compounds

Example

Alkenes

Ethylene

!C#C!

Alkynes

Acetylene

!X (X = F, Cl, Br, I)

Organic halides

Methyl chloride

!OH

Alcohols, phenols

Ethanol, phenol

C!O!C

Ethers

Diethyl ether

Amines

Methylamine

Carboxylic acids

Acetic acid

Amides

Acetanilide

C C

N O C OH O C N

(continues)

52


Table

❚

2.3 (continued )

Some common functional groups

Functional Group

Class of Compounds

Example

Aldehydes

Formaldehyde

Ketones

Methyl ethyl ketone

O C H O C

2.7

Chemical Nomenclature

Although only a limited number of elements exist, the number of compounds that may be formed from those elements is virtually boundless. Given the vast number of molecules that can be made, we require a systematic means of assigning names to chemical compounds. This system should be sufficiently well defined that a person who knows the rules can draw the structure of any compound, given its systematic name. This naming process for molecules is often referred to as chemical nomenclature. We will establish some of the basic premises of this system now and supplement this initial set of rules as needed when we encounter new situations and types of compounds later.

Binary Systems Compounds that contain only two elements are called binary compounds. Fe2O3, for example, is a binary compound. Many such compounds exist and they can be conveniently grouped according to their bonding tendencies. Thus, we encounter slightly different rules when we name binary molecules held together by covalent bonds than when we assign names to ionic compounds. This means it is important that you learn to recognize from its formula whether a compound is ionic or covalent. This will become easier as you get more comfortable working with chemical formulas. A good way to start is to begin to recognize the elements as metals or nonmetals. When two nonmetals combine, they usually form a covalent compound. But when metals and nonmetals combine with one another, they frequently form ionic compounds. It is also handy to learn to recognize common polyatomic ions, such as those listed in Tables 2.5 and 2.6. The presence of these ions is a sign that a compound is ionic.

Naming Covalent Compounds In some cases, a given pair of elements can form compounds in a number of different ways. Nitrogen and oxygen, for example, form NO, N2O, NO2, N2O3, N2O4, and N2O5, all of which are stable enough to observe. So it is critical that our naming system distinguishes these different molecules. To accomplish this, the nomenclature system uses a prefix to specify the number of each element present. The first ten of these prefixes, which arise from the Greek roots for the numbers, are listed in Table 2.4. In a binary compound, the element that appears first in the formula also appears first in the name. The first element retains its full name, whereas the second element

Table

❚

2.4

Greek prefixes for the first ten numbers

Number

Prefix

One

Mono-

Two

Di-

Three

Tri-

Four

Tetra-

Five

Penta-

Six

Hexa-

Seven

Hepta-

Eight

Octa-

Nine

Nona-

Ten

Deca-


53

(a) Dinitrogen oxide, N2O

(d) Dinitrogen trioxide, N2O3

(b) Nitrogen monoxide, NO

(e) Dinitrogen tetroxide, N2O4

(c) Nitrogen dioxide, NO2

(f) Dinitrogen pentoxide, N2O5

Nitrogen forms a number of binary compounds with oxygen.

is described by replacing the ending from its name with the suffix -ide. Both elements will be preceded by a number-designating prefix except that when there is only one atom of the first element, it does not carry the prefix mono-. An example of this procedure is provided in Example Problem 2.5.

E XAM PL E PRO BLEM

2 .5

What are the systematic names of the following compounds? (a) N2O5, (b) PCl3, (c) P4O6

Strategy The first element listed retains its full name and adds a prefix when more than one atom of it is in the compound. The second element will retain only the root of its name, followed by -ide, and it too takes a prefix to indicate the number of atoms.

Solution (a) N2O5: dinitrogen pentoxide, (b) PCl3: phosphorus trichloride (remember: this is not called monophosphorus trichloride), (c) P4O6: tetraphosphorus hexoxide. (The a in hexa- is dropped here to simplify pronunciation.)

Check Your Understanding What are the names of the following compounds? (a) CS2, (b) SF6, (c) Cl2O7

Naming Ionic Compounds The iron chlorides shown earlier in Figure 2.16 are examples of binary ionic compounds. Because ionic compounds must be neutral, the positive and negative charges of the ions must balance each other and only one formula unit is possible. Therefore, once one of the charges is specified in the name, the entire formula is known. The rules of nomenclature specify that the positively charged species, the cation, be specified with enough information about its charge to indicate the complete formula. Unfortunately for novices in chemistry, some groups of the periodic table have only one possible cationic charge, whereas other groups (such as Group 8 that contains iron) have more than one possibility. The most common cations are summarized in Table 2.5. Note that the cations of all Group 1 metals have a charge of 1+ and those of Group 2 have a charge of 2+. Because the nomenclature system was designed by chemists to be used by chemists, it assumes that we know this type of factual information. The nomenclature system gives

54


Table

❚

2.5

Common cations Sodium ion

Na+

Potassium ion

K+

Magnesium ion

Mg2+

Calcium ion

Ca2+

Iron(II) ion

Fe2+

Copper(I) ion

Cu+

Iron(III) ion

Fe3+

Copper(II) ion

Cu2+

Zinc ion

Zn2+

Hydronium ion

H3O+

+

Silver ion

Ag

Ammonium ion

NH4+

Table

❚

2.6

Common anions Halides

F−, Cl−, Br−, I−

Sulfate

SO42−

Nitrate

NO3−

Hydroxide

OH−

Phosphate

PO43−

Cyanide

CN−

Carbonate

2−

Hydrogen carbonate

HCO3−

CO3

no indication of charge for groups that always have the same charge. When we look at the charges in transition metals, however, we note that they often form two cations and some even more than two. When we encounter such a situation, we indicate the charge of the cation in the name using a Roman numeral in parentheses following the elemental name. Thus Fe2+ has the name iron(II), and Fe3+ has the name iron(III). An older system was once used that distinguished these two ions by using a suffix. The smaller charge formed a name ending in -ous (ferrous for Fe2+ in the case of iron whose root, ferr, comes from the Latin word for iron), and the larger charge ended in -ic (ferric for Fe3+). We will not use this system in this book, but you may see these older names on some laboratory chemicals. We’ve already seen some examples of elements that form monatomic anions. The most common are halogens. There are also several polyatomic anions that we will encounter often in this book. Those listed in Table 2.6 are the most common anions that we will see throughout the text. The naming convention for monatomic ions is familiar from the discussion of covalent molecules above: the name of a monatomic anion is the root of the element name with the suffix -ide added. Thus Cl− is chloride, Br− is bromide, and so on. We can now name FeCl2 and FeCl3 as iron(II) chloride and iron(III) chloride, respectively. Many compounds contain polyatomic anions, including those shown in Table 2.6. Most often, the names of these polyatomic ions are memorized rather than being obtained by a systematic nomenclature rule. There is, however, a system for polyatomic anions that contain oxygen and one other element, oxyanions. The base name of the oxyanion is provided by the element that is not oxygen. If there are two possible groupings of the element with oxygen, the one with more oxygen atoms uses the suffix -ate and the one with fewer oxygens uses the suffix -ite. When there are four possible oxyanions, we add a prefix per- to the -ate suffix for the most oxygens and a prefix hypo- to the -ite suffix for the least oxygens. Chlorine is the classic example of an element that forms four oxyanions, whose names are provided in Table 2.7.

Generally, the bonds between the atoms within a polyatomic ion are covalent.

Table

❚

2.7

Oxyanions of chlorine ClO−

Hypochlorite

ClO2−

Chlorite

ClO3−

Chlorate

ClO4−

Perchlorate


55

Once we know how to name both of the ions, an ionic compound is named simply by combining the two names. The cation is listed first in the formula unit and in the name. Example Problem 2.6 provides some examples of the way to determine the name of an ionic compound.

E XAM PL E PRO BLEM

2 .6

Determine the name of the following ionic compounds: (a) Fe2O3, (b) Na2O, (c) Ca(NO3)2

Strategy We must determine the names of the constituent ions first. The anions will provide a hint about the cation charges if we need it.

Solution (a) Fe2O3: As noted in Table 2.6, oxygen is always a 2− ion in these compounds, so there is a total charge of 6− on the three oxide ions in the formula unit. Therefore, the two iron ions must have a total charge of 6+, requiring 3+ from each iron. So the name is iron(III) oxide. (b) Na2O: Sodium from Group 1 always has a 1+ charge and oxygen always 2−. Therefore, the name is sodium oxide. No Roman numeral is needed for sodium because it has only one common ionic charge. (c) Ca(NO3)2: The calcium ion is in Group 2 and always carries a charge of 2+. NO3− is a common polyatomic anion called nitrate. The name is calcium nitrate.

Check Your Understanding Name each of the following ionic compounds: (a) CuSO4, (b) Ag3PO4, (c) V2O5 Occasionally throughout this text, we will encounter new classes of chemical compounds that will need more rules to determine their names. We will introduce these nomenclature systems when necessary.

INSIGHT INTO

2.8

Free radicals are reactive wherever they form. Physiological processes including aging have been tied to the presence of free radicals, so their importance is wide ranging.

56

Polyethylene

We close this chapter by taking a closer look at polyethylene, which is probably the most common polymer in the world. The U.S. domestic production of all forms of polyethylene in 2007 totaled more than 39 billion pounds. And in its various forms, polyethylene can be found in items ranging from plastic grocery bags to children’s toys, automobile gas tanks, and even bulletproof vests. Let’s start at the beginning, with a look at the way polyethylene is produced. In Section 2.4, we pointed out that polyethylene is built up from a monomer called ethylene and that the molecular formula for ethylene is C2H4. We also noted that the formula for polyethylene itself is often written as ![CH2CH2]n!. To begin to see how the monomers can be combined to form the polymer, we need to look at the structural formula for ethylene, which is shown on page 57. From our discussion of line structures, we know that each carbon atom in a molecule always forms a total of four covalent bonds. For the carbon atoms in ethylene, two of those bonds are to hydrogen atoms, and the other two make up a double bond between the two carbon atoms. To link the monomers together and form the polymer, then, we will need to convert this double bond into a single bond, leaving each carbon with one bond available for linking to the next monomer in the chain. The polymerization reaction begins with the addition of a small amount of an initiator molecule


that easily breaks down, producing highly reactive fragments called free radicals. (The free radical is denoted as ‘R?’ below, where the ‘?’ represents an unshared electron.) One of these free radicals attaches itself to a single ethylene molecule, opening its double bond and leaving one end unbonded. H R?

H C

C

H

R H

H

H

C

C?

H

H

The unbonded end of the ethylene now takes on the role of the free radical, attacking a second ethylene monomer and attaching it to the growing chain. As long as the number of available ethylene monomers remains large, the polymer can continue to grow in this fashion.

R

H

H

C

C?

H

H

H

H C

C

H

R H

H

H

H

H

C

C

C

C?

H

H

H

H

Linking together thousands of monomers in this way would generate a single polyethylene molecule, whose structural formula would look like that below, only extending over a much greater chain length. H

H

H

H

H

H

H

H

H

H

C

C

C

C

C

C

C

C

C

C

H

H

H

H

H

H

H

H

H

H

Eventually, the free radical end of the growing chain will meet up with another free radical, either from an initiator molecule or from another growing chain. When this happens, the chain will stop growing. When the monomer units grow end to end, as shown above, the result is known as linear polyethylene because all of the carbon atoms lie along a more or less straight backbone. It is also possible to grow polyethylene under conditions that lead to branched chains, in which some of the hydrogen atoms along the backbone are themselves replaced by polyethylene-like chains. The contrasting forms of the linear and branched molecules are shown in Figure 2.17. Although linear polyethylene is conceptually simpler, branched polyethylene is actually cheaper and easier to produce. Perhaps surprisingly, the linear and branched forms of polyethylene are actually polymers with significantly different macroscopic properties. When arranged side by side, the long, linear polyethylene molecules can pack together very tightly, producing a relatively dense plastic. So linear polyethylene is also known as high-density polyethylene, or HDPE. It is a strong and hard material, used in bottles, kitchenware, and as a structural plastic in many children’s toys. Because their chains cannot be stretched out straight, branched polyethylene molecules cannot be packed together as closely as those of linear polyethylene. This looser packing of the molecules produces a plastic with a much lower density, so branched polyethylene is usually referred to as LDPE, for low-density polyethylene. LDPE is commonly used in applications where relatively little strength is needed, such as plastic films, sandwich bags, and squeeze bottles. Recent advances in polymerization technology have made it possible to grow extremely long linear polyethylene chains, stretching to hundreds of thousands of monomer units. Because the individual molecules are relatively large and heavy, this has been called “ultra-high molecular weight polyethylene,” or UHMWPE. The very

2.8

Polyethylene

57

Figure 2.17 ❚ Differences between linear and branched polyethylene are illustrated. The left-hand panel shows linear, or high-density polyethylene, and the right-hand panel shows the branched, or low-density form. In each case, the upper diagram illustrates the molecular structure for part of a polymer chain. The lower diagrams show the way that polymer chains would pack together to form the solid plastic. (Hydrogen atoms are omitted in these drawings for clarity.) Branched chains cannot approach one another as closely, so the resulting material has a much lower density.

Linear polyethylene (high-density)

Branched polyethylene (low-density)

long chains are so strong that this material is replacing Kevlar® (another polymer) as the standard filling for bulletproof vests. UHMWPE can also be formed into large sheets, and these have been used as ice substitutes for skating rinks. This quick glimpse at some of the many forms and uses for polyethylene shows how the observable properties of polymers are closely linked to the chemical structure of the individual molecules. It also illustrates how the interplay between chemistry and engineering can lead to polymers that are selectively produced to meet certain design specifications.

FOCU S O N

PRO BLEM S O LVI NG

Question Boron is widely used in the production of enamels and glasses. Naturally occurring boron has an average atomic mass of 10.811 amu. If the only isotopes present are 10B and 11B, describe how you would determine their relative abundances. Include in your description any information that you would need to look up.

Strategy This problem is the inverse of those in this chapter in which we used isotopic abundances to find atomic masses. Here we must work from the atomic mass to find the isotopic abundances. Since there are two isotopes, we have two unknowns. So we will need to write two equations that relate the percentages or fractions of 10B and 11B. We would also need to know the mass of each isotope, but presumably we could look those up.

Solution The abundances of the two isotopes must add to 100%. This gives us a first equation: (Fraction of 10B) + (fraction of 11B) = 1.00 The fact that the average atomic mass is 10.811 provides another equation: (Fraction of 10B) × (mass of 10B) + (fraction of 11B) × (mass of 11B) = 10.811 amu

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Assuming that we are able to find values for the masses of the two isotopes, we now have two equations in two unknowns, so the problem can be solved. (We could get a rough approximation by assuming that 10B has a mass of 10 amu and 11B has a mass of 11 amu, but that assumption would not lead to a very accurate result.)

S U M M ARY The widely taught description of an atom as a massive, positively charged nucleus surrounded by lighter, fast moving electrons is based on the results of many ingenious experiments. Several details of this model are important in building our knowledge of chemistry. The number of protons in a nucleus determines the chemical identity of an atom, and must be equal to the number of electrons so that the atom is electrically neutral. If the numbers of protons and electrons are not the same, the result is a charged particle, which is called an ion. Because of their electric charge, ions behave quite differently from neutral atoms. For example, the sodium in our diet is invariably sodium ions, not sodium atoms. Electric charge plays a central role in determining the structure of both atoms and molecules. The attraction of oppositely charged particles and the repulsion of like charged particles, as described by Coulomb’s law, is central to our understanding of the way atoms stick together, that is, chemical bonding. The various types of chemical bonds, including covalent, ionic, and metallic, can all be understood in terms of the interactions between negative and positive charges. Once we establish that atoms will bond together to form chemical compounds, we are faced with the burden of summarizing a vast number of compounds. Many classification schemes

have been developed over the years to assist us in organizing data on a wide range of chemical substances. The periodic table is the most common and important device for such a purpose. It succinctly summarizes many properties of the elements, especially their chemical tendencies. The periodic table also helps us learn chemistry by providing a template for trends. Thus we can remember that metallic elements are found toward the left and bottom of the table while nonmetals are found toward the upper right. Without the periodic table, remembering which elements are in each category would be much more challenging. Other categorizations also help organize our study of chemistry. In some cases, broad categories such as organic versus inorganic chemistry are helpful. For example, chemical nomenclature, the system we use to name compounds, is different for these two branches of chemistry. The way that we impart information about molecules symbolically also varies. In organic chemistry, because carbon is involved in all of the molecules, we use the fact that carbon always forms four chemical bonds to simplify the depiction of molecules as line structures. For inorganic chemistry, in which we often encounter binary chemical compounds, we can devise a fairly broad system of nomenclature based on a relatively small number of rules. Learning this system is an important step in chemical communication.

K EY TE RM S actinides (2.5)

formula unit (2.4)

main group elements (2.5)

addition reaction (2.6)

free radicals (2.8)

mass number (2.2)

alkali metals (2.5)

functional group (2.6)

metal (2.5)

alkaline earths (2.5)

groups (2.5)

metallic bonding (2.4)

anion (2.3)

halogens (2.5)

metalloid (2.5)

atomic mass unit (2.2)

HDPE (2.8)

molecular formula (2.4)

atomic number (2.2)

hydrates (2.4)

molecule (2.4)

binary compound (2.7)

hydrocarbons (2.6)

monomer (2.1)

cation (2.3)

inorganic chemistry (2.6)

neutrons (2.2)

chemical bond (2.4)

ion (2.3)

noble gases (2.5)

chemical compound (2.4)

ionic bonding (2.4)

nonmetal (2.5)

chemical formula (2.4)

isotopes (2.2)

nucleus (2.2)

chemical nomenclature (2.7)

isotopic abundance (2.2)

organic chemistry (2.6)

Coulomb’s law (2.3)

lanthanides (2.5)

oxyanions (2.7)

covalent bonds (2.4)

lattice (2.4)

periodic law (2.5)

electrons (2.2)

LDPE (2.8)

periodicity (2.5)

empirical formula (2.4)

line structure (2.6)

periods (2.5)

Key Terms

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polyatomic ions (2.7)

representative elements (2.5)

transition metals (2.5)

polymer backbone (2.1)

semimetals (2.5)

UHMWPE (2.8)

protons (2.2)

P RO B L E M S AND E XE RCISE S ■

denotes problems assignable in OWL.

2.18

INSIGHT INTO Polymers 2.1 Define the terms polymer and monomer in your own words. 2.2 How do polymers compare to their respective monomers?

The element gallium, used in gallium arsenide semiconductors, has an atomic mass of 69.72 amu. There are only two isotopes of gallium, 69Ga with a mass of 68.9257 amu and 71Ga with a mass of 70.9249 amu. What are the isotopic abundances of gallium? ■

2.3 Look around you and identify several objects that you think are probably made from polymers.


2.4 Which one element forms the backbone of nearly all common polymers? Which other elements are also found in common household polymers? 2.5 The fact that a polymer’s physical properties depend on its atomic composition is very important in making these materials so useful. Why do you think this would be so? 2.6 Use the web to research the amount of PVC polymer produced annually in the United States. What are the three most common uses of this polymer? 2.7 Use the web to research the amount of polyethylene produced annually in the United States. What are the three most common uses of this polymer? 2.19

Atomic Structure and Mass 2.8 In a typical illustration of the atom such as Figure 2.3, which features lead to misunderstandings about the structure of atoms? Which ones give important insight? 2.9 Why is the number of protons called the atomic number? 2.10

Gallium melts just above room temperature. ■ The atomic mass of copper is 63.55 amu. There are only two isotopes of copper, 63Cu with a mass of 62.93 amu and 65 Cu with a mass of 64.93 amu. What is the percentage abundance of each of these two isotopes?

2.20 The following table presents the abundances and masses of the isotopes of zinc. What is the atomic mass of zinc?

Which isotope in each pair contains more neutrons? (a) chlorine-35 or sulfur-33, (b) fluorine-19 or neon-19, (c) copper-63 or zinc-65, (d) iodine-126 or tellurium-127

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Isotope

Abundance

Mass

64

48.6%

63.9291 amu

Write the complete atomic symbol for each of the following isotopes: (a) carbon-13, (b) phosphorus-31, (c) sodium-23, (d) boron-10

66

27.9%

65.9260 amu

67

4.10%

66.9271 amu

How many electrons, protons, and neutrons are there in each of the following atoms? (a) magnesium-24, 24Mg, (b) tin-119, 119Sn, (c) thorium-232, 232Th, (d) carbon-13, 13C, (e) copper-63, 63Cu, (f) bismuth-205, 205Bi

68

18.8%

67.9249 amu

70

0.60%

69.9253 amu

Zn

2.11 Define the term isotope. 2.12

2.13

2.14

Zn Zn

■

Consider the following nuclear symbols. How many protons, neutrons, and electrons does each element have? What elements do R, T, and X represent? ■

(a)

30 14R

(b)

89 39T

(c)

133 55X

2.15 Mercury is 16.716 times more massive than carbon-12. What is the atomic mass of mercury? Remember to express your answer with the correct number of significant figures. 2.16 How can an element have an atomic mass that is not an integer? 2.17 Explain the concept of a “weighted” average in your own words.

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Zn

■


Zn

2.21

Naturally occurring uranium consists of two isotopes, whose masses and abundances are shown below.

■

Isotope

Abundance

Mass

235

0.720%

235.044 amu

238

99.275%

238.051 amu

U U

Only 235U can be used as fuel in a nuclear reactor, so uranium for use in the nuclear industry must be enriched in this isotope. If a sample of enriched uranium has an average atomic mass of 235.684 amu, what percentage of 235U is present?

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Ions 2.22 Define the term ion in your own words. 2.23 What is the difference between cations and anions? 2.24

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2.25

■

Provide the symbol of the following monatomic ions, given the number of protons and electrons in each: (a) 8 protons, 10 electrons, (b) 20 protons, 18 electrons, (c) 53 protons, 54 electrons, (d) 26 protons, 24 electrons

How many protons and electrons are in each of the following ions? (a) Na+, (b) Al3+, (c) S2−, (d) Br−

2.26 Identify each of the following species as an anion, a cation, or a molecule: (a) CO32−, (b) CO2, (c) NH4+, (d) N3−, (e) CH3COO− 2.27

■ Write

the atomic symbol for the element whose ion has a 2– charge, has 20 more neutrons than electrons, and has a mass number of 126.

2.28 In what region of the periodic table are you likely to find elements that form more than one stable ion? 2.29

Give the symbol, including the correct charge, for each of the following ions. (a) barium ion, (b) titanium(IV) ion, (c) phosphate ion, (d) hydrogen carbonate ion, (e) sulfide ion, (f) perchlorate ion, (g) cobalt(II) ion, (h) sulfate ion ■

2.30 An engineer is designing a water softening unit based on ion exchange. Use the web to learn what ions typically are “exchanged” in such a system. Given that the ion exchanger cannot build up a large positive charge, what can you conclude about the relative numbers of the various ions involved? 2.31 Use the web to find a catalyst for a polymerization reaction that uses an ion. What are the apparent advantages of using this catalyst for creating the polymer? 2.32 Using Coulomb’s law, explain how the difference between attractive and repulsive interactions between ions is expressed mathematically. Compounds and Chemical Bonds 2.33 How many atoms of each element are represented in the formula Ba(OH)2? 2.34 Which of the following formulas contains the most hydrogen atoms? C2H6, (NH4)2CO3, H2SO4, or Fe(OH)3. 2.35 In general, how are electrons involved in chemical bonding? 2.36 What is the difference between an ionic bond and a covalent bond? 2.37 When talking about the formula of an ionic compound, why do we typically refer to a formula unit rather than a molecule? 2.38 Which formula below is correct for an ionic compound? What is incorrect about each of the others? (a) AlO3/2, (b) Ca2Br4, (c) Mg(PO4)3/2, (d) BaCO3 2.39 Explain the differences between ionic and metallic bonding.

2.42 Explain the difference between a molecular formula and an empirical formula. 2.43 Why are empirical formulas preferred for describing polymer molecules? 2.44 The molecular formula for the ethylene monomer is C2H4. What is its empirical formula? 2.45 Polybutadiene is a synthetic elastomer, or rubber. The corresponding monomer is butadiene, which has the molecular formula C4H6. What is the empirical formula of butadiene? The Periodic Table 2.46 What distinguished the work of Mendeleev that caused scientists to accept his concept of the periodic table when others before him were not believed? 2.47 How does the periodic table help to make the study of chemistry more systematic? 2.48 What is a period in the periodic table? From what does it derive its name? 2.49 How do binary compounds with hydrogen illustrate the concept of periodicity? 2.50

Of the following elements, which two would you expect to exhibit the greatest similarity in physical and chemical properties? Cl, P, S, Se, Ti. Explain your choice.

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2.51 Name the group to which each of the following elements belongs: (a) K, (b) Mg, (c) Ar, (d) Br 2.52 What are some of the physical properties that distinguish metals from nonmetals? 2.53

■ Identify the area of the periodic table in which you would expect to find each of the following types of elements. (a) a metal, (b) a nonmetal, (c) a metalloid

2.54 Why are nonmetals important even though they account for only a very small fraction of the elements in the periodic table? 2.55 What is a metalloid? 2.56 A materials engineer has filed for a patent for a new alloy to be used in golf club heads. The composition by mass ranges from 25 to 31% manganese, 6.3 to 7.8% aluminum, 0.65 to 0.85% carbon, and 5.5 to 9.0% chromium, with the remainder being iron. What are the maximum and minimum percentages of iron possible in this alloy? Use Figure 2.12 to make a prediction about how the density of this alloy would compare with that of iron; justify your prediction. 2.57

■ Classify the following elements as metals, metalloids, or nonmetals: (a) Si, (b) Zn, (c) B, (d) N, (e) K, (f) S

2.58 A materials engineer wants to make a new material by taking pure silicon and replacing some fraction of the silicon atoms with other atoms that have similar chemical properties. Based on the periodic table, what elements probably should be tried first?

2.40 Conduction of electricity usually involves the movement of electrons. Based on the concept of metallic bonding, explain why metals are good conductors of electricity.


2.41 Describe how a covalently bonded molecule is different from compounds that are either ionic or metallic.

2.60 Calcium and fluorine combine to produce ionic calcium fluoride, CaF2. Use the periodic table to predict at least

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2.59 The chemistry of main group elements is generally simpler than that of transition metals. Why is this so?


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are due to the interaction between ethanol and GABA. Write the correct molecular formula for this compound.

two other compounds that you would expect to have structures similar to that of CaF2. 2.61 What is meant by the phrase organic chemistry? 2.62 Based on what you’ve learned in this chapter, would you classify the chemistry of polymers as organic or inorganic? Why? 2.63 What is a functional group? How does the concept of the functional group help to make the study of organic chemistry more systematic? 2.64 The molecule shown below is responsible for the smell of popcorn. Write the correct molecular formula for this compound.

OH

H2N O

2.68 The figure below shows the structure of the adrenaline molecule. Write the correct molecular formula for this substance.

OH OH

H N

N OH O Chemical Nomenclature 2.65

Not all polymers are formed by simply linking identical monomers together. Polyurethane, for example, can be formed by reacting the two compounds shown below with one another. Write molecular and empirical formulas for each of these two substances. ■

2.69 Why are there different rules for naming covalent and ionic binary compounds? 2.70 Which binary combinations of elements are most likely to give ionic substances? 2.71

■ Name the following covalent compounds: (a) N 2 O 5 , (b) S2Cl2, (c) NBr3, (d) P4O10

2.72

■ Give the formula for each of the following compounds: (a) sulfur dichloride, (b) dinitrogen pentaoxide, (c) silicon tetrachloride, (d) diboron trioxide (commonly called boric oxide)

2.73

■ Write the molecular formula for each of the following covalent compounds: (a) sulfur hexafluoride, (b) bromine pentafluoride, (c) disulfur dichloride, (d) tetrasulfur tetranitride

2.74

■ Name each of the following ionic compounds: (a) K2S, (b) CoSO4, (c) (NH4)3PO4, (d) Ca(ClO)2

2.75

■ Name each of the following compounds: (a) MgCl 2 , (b) Fe(NO3)2, (c) Na2SO4, (d) Ca(OH)2, (e) FeSO4

2.76

■ Give the formula for each of the following ionic compounds: (a) ammonium carbonate, (b) calcium iodide, (c) copper(II) bromide, (d) aluminum phosphate, (e) silver(I) acetate

2.77

■ Name the following compounds: (a) PCl5, (b) Na2SO4, (c) Ca3N2, (d) Fe(NO3)3, (e) SO2, (f ) Br2O5

O C N

N

HO

C

OH

O 2.66

The compound shown below forms an amorphous solid (a glass) at room temperature and has been used as a medium for storing information holographically. Write the correct molecular formula for this molecule.

■

O

N C

O

INSIGHT INTO Polyethylene 2.78 What is a free radical? How are free radicals important in the formation of polyethylene?

N

2.67 The accompanying figure shows the structure gammaaminobutanoic acid, or GABA. This molecule is a neurotransmitter. Some of the effects of alcohol consumption

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2.79 How do molecules of low-density polyethylene and highdensity polyethylene differ? How do these molecular scale differences explain the differences in the macroscopic properties of these materials? 2.80 Why do you think an initiator molecule is needed to induce the polymerization of ethylene? 2.81 Use the web to determine the amount of low-density polyethylene and high-density polyethylene produced annually

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in the United States. Which uses predominate in the applications of these two materials?

FOCUS ON PROBLEM SOLVING EXERCISES 2.82 Describe how you can identify the isotope, X, in this puzzle. The nucleus contains one more neutron than proton, and the mass number is nine times larger than the charge on the ion X3+. 2.83 Many transition metals produce more than one ion. For example, iron has ions with charges of 2+ and 3+ that are both common. How could you use the compounds of a transition metal with oxygen to determine the charge of the metal ion? Use iron as your example. 2.84 Naturally occurring europium has an average atomic mass of 151.965 amu. If the only isotopes of europium present are 151Eu and 153Eu, describe how you would determine the relative abundance of the two isotopes. Include in your description any information that would need to be looked up. 2.85

■ Strontium has four stable isotopes. Strontium-84 has a very low natural abundance, but 86Sr, 87Sr, and 88Sr are all reasonably abundant. Knowing that the atomic mass of strontium is 87.62, which of the more abundant isotopes predominates?

2.86 A candy manufacturer makes chocolate covered cherries. Although all of the products look roughly the same, 3% of them are missing the cherry. The mass of the candy with a cherry is 18.5 g; those missing the cherry weigh only 6.4 g. (a) How would you compute the average mass of a box of

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100 of these chocolate covered cherries from this manufacturer? (b) How is this question analogous to the determination of atomic masses? 2.87 Two common oxides of iron are FeO and Fe2O3. Based on this information, tell how you would predict two common compounds of iron and chlorine. Cumulative Problems 2.88 Use a molecular level description to distinguish between LDPE and HDPE. 2.89 Engineers who design bicycle frames are familiar with the information in Table 1.3, which lists the densities of aluminum (2.699 g/cm 3), steel (7.87 g/cm 3), and titanium(4.507 g/cm3). How does this information compare with Figure 2.12 and what would it suggest for changes in this figure if more shades were used for the density colorcoding? (Iron is the principle component of steel.) 2.90 Use the web to look up the density of different forms of steel, such as stainless steel or magnetic steel, and discuss whether or not the differences in the densities follow what might be predicted by looking at the periodic properties of elements. 2.91 LDPE has a density in the range of 0.915−0.935 g/cm3, and HDPE has a density in the range of 0.940−0.965 g/cm3. You receive a small disk, 2.0 cm high with a 6.0-cm diameter, from a manufacturer of polyethylene, but its label is missing. You measure the mass of the disk and find that it is 53.8 g. Is the material HDPE or LDPE?


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3 OUTLINE 3.1

INSIGHT INTO Explosions

3.2

Chemical Formulas and Equations

3.3

Aqueous Solutions and Net Ionic Equations

3.4

Interpreting Equations and the Mole

3.5

Calculations Using Moles and Molar Masses

3.6

INSIGHT INTO Explosives and Green Chemistry

Molecules, Moles, and Chemical Equations

Explosions can be very large, as in the photo on the left, or quite small, as in the laboratory demonstrations in the photo on the right. But all chemical explosions are very rapid reactions that release substantial amounts of energy in a short time. Left: U.S. Department of Energy; right: © Cengage Learning/Charles D. Winters

T Online homework for this chapter may be assigned in OWL.

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oday, roughly 20 million chemical compounds are known. More are being discovered or created as you read this. Although chemical nomenclature provides a systematic way to facilitate discussion of this tremendous diversity of substances, that is just a small step toward understanding chemistry. Much of the science of chemistry is aimed at understanding the reactions that produce or consume chemical substances. You may not realize it, but chemical reactions are central to many of the technologies that you use. Chemical reactions in batteries produce electricity for your iPhone® and laptop computer, and the burning of hydrocarbons powers your car. Other reactions, such as those designed to synthesize the light-filtering compounds used in LCD computer displays, play less obvious roles in technology. Over the course of this text we will explore many aspects of chemical reactions, and we will also look at a number of examples to show how chemical reactions can be put to good use through clever engineering. We’ll begin our discussion by considering a class of reactions that can be either useful or destructive but are always dramatic: explosions.

Chapter Objectives After mastering this chapter, you should be able to ❚ list at least three characteristics of explosive chemical reactions. ❚ explain balancing a chemical equation as an application of the law of conservation of mass. ❚ list at least three quantities that must be conserved in chemical reactions. ❚ write balanced chemical equations for simple reactions, given either an unbalanced equation or a verbal description. ❚ explain the concept of a mole in your own words. ❚ interpret chemical equations in terms of both moles and molecules. ❚ interconvert between mass, number of molecules, and number of moles. ❚ determine a chemical formula from elemental analysis (i.e., from % composition). ❚ define the concentration of a solution and calculate the molarity of solutions from appropriate data. ❚ calculate the molarity of solutions prepared by dilution or calculate the quantities needed to carry out a dilution to prepare a solution of a specified concentration. ❚ distinguish between electrolytes and nonelectrolytes and explain how their solutions differ. ❚ describe the species expected to be present (ions, molecules, etc.) in various simple solutions. ❚ recognize common strong acids and bases. ❚ write molecular and ionic equations for acid–base neutralization reactions.

INSIGHT IN TO

3.1

Explosions

From the imposing sound to the frightening potential for destruction, nothing about an explosion is easily ignored. Whether it is a controlled detonation to bring down an old building, a violent industrial accident, or a deliberate bombing, an explosion always commands our attention. If we look at an explosion scientifically, we will usually find a chemical reaction at its heart. Many different chemical reactions can produce explosions. Some are strikingly simple, at least in terms of the substances involved: a mixture of hydrogen and oxygen gases will explode with a distinctive popping sound if ignited with a spark or a flame. Others can be much more complex, involving several large molecules. What common characteristics can we identify that might contribute to a reaction becoming explosive? The most fundamental requirement for an explosive reaction is that it must release a large amount of energy. This release of energy is responsible for what might be described as the “force” of the explosion. The actual amount of energy released in a particular detonation will depend on both the identity and the quantity of the explosive used. In many cases, the energy release occurs as a fairly complex compound (the explosive) decomposes to form a collection of smaller, simpler molecules. The fact that explosives can decompose in a way that releases large amounts of energy means that these compounds are not very stable, and this raises obvious safety concerns for the storage or handling of what have come to be called energetic materials. An explosive is most useful if it can be handled safely and then detonated readily on demand, but these objectives are not always mutually compatible. The same properties that lead to safe handling can also make detonation more difficult. Another common trait of explosive reactions is that they happen very fast; this is what makes an explosion much more dramatic and destructive than a slow burning fire

We will study the role of energy in chemical reactions in detail in Chapter 9.

We will examine the factors that govern the rates of chemical reactions in Chapter 11.

3.1

Explosions

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Figure 3.1 ❚ Alfred Nobel’s invention of dynamite provided a powerful explosive in solid form. The actual explosive in dynamite is nitroglycerin, whose molecular structure is shown. Nitroglycerin is a liquid under ordinary conditions, so its use as a commercial explosive was extremely hazardous.

H2C

O

NO2

HC

O

NO2

H 2C

O

NO2

Nitroglycerin

that releases the same amount of energy. Careful examination of the progress of explosive chemical reactions reveals that they accelerate as they proceed. As a result, all of the available explosive is consumed in a very short period of time. As that happens, the energy from the explosion is also released on that same short timescale. You may know from physics that power is defined as energy per unit time. So as the time required for the explosion to occur becomes shorter, the power of the explosion will increase. Modern explosives are generally solids. But that was not always true; liquid nitroglycerin was one of the first widely used explosives. Liquids are much harder to transport and handle than solids, though, and early use of nitroglycerin was plagued by many accidents due to unintentional detonation. The hazards surrounding the use of liquid nitroglycerin prompted Alfred Nobel to develop dynamite, in which nitroglycerin is mixed with an inert binder to produce a solid material that is much easier and safer to handle (Figure 3.1). In the chemical reactions associated with explosions, the solid explosive is converted predominantly into gases. This is no coincidence; the generation of gases is actually very important to the development of the explosion. When a chemical reaction transforms a solid explosive into a large number of gaseous molecules, those gases will initially occupy a volume similar to that of the solid. You probably know that gases have much lower densities than solids. So the gaseous products tend to expand rapidly to reach their normal densities. This expansion typically results in the formation of a shock wave, which helps to deliver the blow of the explosion to its surroundings as shown in Figure 3.2. When considering chemical reactions—including explosions—it is often handy to think in terms of the microscopic, macroscopic, and symbolic perspectives introduced in Section 1.2. Now we will develop these ideas further in the specific context of reactions and explosions. Figure 3.2 ❚ The destructive force of an explosion is due in part to the expansion of gases produced in the reaction. This molecular scale diagram illustrates the effect. In a solid explosive, relatively large molecules are packed very closely together. When the reaction occurs, the products are usually smaller gas molecules. These gases are initially formed in a volume similar to that of the solid. Because densities are typically hundreds of times smaller in gases than in solids, the products are very highly compressed. The gas must expand to lower its density, and this produces a shock wave.

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Tightly packed molecules in solid explosive

Detonation

Chapter 3 Molecules, Moles, and Chemical Equations

Formation of smaller gas molecules in products, initially highly compressed

Expansion of gas

Gas molecules continue to expand toward normal density.

3.2

Chemical Formulas and Equations

Although the properties and uses of different explosives vary considerably, all involve chemical reactions. A realistic description of the chemistry of virtually any explosion would require us to consider many individual reactions. To discuss the science behind an explosion, we must be able to describe a chemical reaction concisely. In Chapter 2, we’ve already seen how chemical formulas provide a concise way to represent chemical compounds, and we will now describe how chemical equations build on this to accomplish the same goal for chemical reactions.

Writing Chemical Equations Chemical equations are designed to represent the transformation of one or more chemical species into new substances. To ensure that their meanings are clear, we follow a set of conventions when writing chemical equations. Each chemical equation has two sides, and we usually envision the reaction as proceeding from left to right. The original materials are called the reactants and they appear on the left-hand side of the equation. The compounds that are formed from the reaction are called products and appear on the right-hand side of the equation. An arrow is used to represent the changes that occur during the reaction. Thus we can write a completely generic chemical equation: Reactants : Products This would usually be read as “reactants go to products” or “reactants give products.” We use chemical formulas to identify the specific reactants and products. The physical states of the compounds are often designated; (s) indicates a solid, (,) a liquid, (g) a gas, and (aq) a substance dissolved in water. We can write the reaction of hydrogen with oxygen, which is often done as a lecture demonstration, as a simple example.

When a substance dissolves in water, the result is called an aqueous solution, and the “aq” label is short for aqueous.

2 H2(g) + O2(g) : 2 H2O(g) All substances involved in this equation are gases, as indicated by the “(g)” notation. One characteristic of this reaction is that it will only occur at high temperatures. As seen in the photos in Figure 3.3, the mixture of H2 and O2 is stable unless ignited. If our written equations are going to describe fully the chemical processes they represent, then we need ways to indicate such conditions in addition to identifying the chemicals involved. Generally, this is done by placing a symbol over the arrow. Here we will introduce two such symbols: one for heat and one for light. A reaction that requires heat (or high temperatures) is indicated with a Δ (the Greek letter delta)

Photos: © Cengage Learning/ Charles D. Winters

Figure 3.3 ❚ A balloon filled with

2 H2(g) ⫹ O2(g)

2 H2O(g)

a mixture of H2 and O2 explodes in a classroom demonstration. The gas mixture in the balloon is stable until ignited by the flame from a candle. Once ignited, the balloon explodes in a ball of flame as the reaction takes place. The accompanying microscopic scale illustration shows the molecular species present before and after reaction.

⫹ Reactants

Products

3.2 Chemical Formulas and Equations

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above the arrow, and one that is initiated by light energy is designated with an hn. We show two examples of these symbols below. Δ 2 H O(g) 2 H2(g) + O2(g) 9: 2

The symbol hn is used because it represents the energy of the light, as we will learn in Chapter 6.

H2(g) + Cl2(g)

hn 9:

2 HCl(g)

The first equation is a more complete description of the reaction between hydrogen and oxygen. In the second equation, light energy is used to initiate the reaction of hydrogen with chlorine, producing hydrogen chloride. Such a light-induced process is called a photochemical reaction.

Balancing Chemical Equations

If we also consider nuclear reactions, we must acknowledge that matter and energy can be interconverted. Nuclear chemistry is discussed in Chapter 14.

Figure 3.4 ❚ The chemical reaction for the burning of methane (CH4) in oxygen illustrates the concept of atom balance for chemical equations. The equation is shown first in symbolic form, with each compound represented by its chemical formula. The individual molecules are pictured in the center row and then are broken down into their constituent atoms. The number of atoms of each element is the same on the left- and right-hand sides.

If you look carefully at the reactions we’ve been writing, you’ll notice more than just the formulas and the states of the substances in the equation. Numerical information about the relative amounts of the substances involved is also given. For example, we included the number “2” in front of both the H2(g) and H2O(g) in the equation for the reaction between hydrogen and oxygen. Why are these numbers there, and what do they mean? If you’ve studied chemistry before in high school, you probably remember that these coefficients are often needed to balance a chemical equation. The underlying premise of the chemical equation is that it is a written representation of a chemical reaction. So any reasonable representation must be consistent with all of our observations of the actual reaction. One of the most fundamental laws of nature is the law of conservation of matter: matter is neither created nor destroyed. If we specifically exclude nuclear reactions from our consideration, this law can be phrased more specifically. Atoms of one element are neither created nor destroyed in a chemical reaction. A chemical reaction simply rearranges the atoms present into new compounds. In its written representation of nature, therefore, the chemical equation must not “create or destroy” atoms. To uphold this condition, we must have the same number of atoms of each element on both sides of the chemical equation (see Figure 3.4). An equation that does not meet this condition cannot accurately represent the observed chemical reaction and is said to be unbalanced. In many cases, the most efficient way to balance the equations is “by inspection,” which really means by trial and error. Even when this exploratory method is used, however, some systematic strategies can help make the balancing process easier. In all cases, the way we obtain balance is to introduce numbers in front of appropriate

+

CH4

Equation

CO2

+

2 H2O

Molecules

Atoms One carbon atom

Four hydrogen atoms Reactants

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2 O2


Four oxygen atoms

One carbon atom

Four oxygen atoms Products

Four hydrogen atoms

formulas in the equation. Chemists use the term stoichiometry to refer to the various quantitative relationships between the amounts of reactants and products in a chemical reaction. So the numbers used to balance a chemical equation are called stoichiometric coefficients. The stoichiometric coefficient multiplies the number of atoms of each element in the formula unit of the compound that it precedes. Example Problem 3.1 shows one strategy for balancing a common class of reactions, the combustion of hydrocarbon fuels. Although these combustion reactions obviously release energy, in most instances they are carried out under conditions at which the reaction is not fast enough to produce a full-fledged explosion.


3 .1

Propane, C3H8, is used as a fuel for gas barbecue grills, where it burns in a controlled fashion. But if a mixture of propane and air is ignited in a closed space, like a gas pipeline, an explosion can easily result. In either of these cases, the propane combines with oxygen, O2, to form carbon dioxide and water. Write a balanced chemical equation describing this reaction.

Strategy This problem requires two steps. First we must read the problem and determine which substances are reactants and which are products. This will allow us to write an unbalanced “skeleton” equation. Then we must proceed to balance the equation, making sure that the same number of atoms of each element appears on both sides.

Solution Step 1: The problem notes that propane burns or explodes by combining with oxygen. The reactants are therefore C3H8 and O2. We also are told that the products are CO2 and H2O. So we have enough information to write the unbalanced equation. (Here we will write “blanks” in front of each formula to emphasize that we still need to determine the coefficients.) –––– C3H8 + –––– O2

––––CO2 + –––– H2O

Step 2: Balancing an equation like this is aided by making some observations. In this case, both carbon and hydrogen appear in only one place on each side of the equation. Our first steps will be to balance these two elements, because there will be no other way to adjust them. Let’s begin with carbon. (This choice is arbitrary.) To achieve the required balance, we can use stoichiometric coefficients of one for C3H8 and three for CO2, giving three carbon atoms on each side of the equation. At this point, we have 1 C3H8 + –––– O2

3 CO2 + –––– H2O

(The coefficient of one in front of the propane would normally be omitted, but is written explicitly here for emphasis.) Next we will balance hydrogen. The propane already has its coefficient set from our work on carbon, so it dictates that there are


69

eight H atoms to be accounted for by water on the product side of the equation. We need to insert a coefficient of four to achieve this balance. 1 C3H8 + –––– O2

3 CO2

+

4 H2O

This leaves us with only oxygen to balance. We have saved this for last because it appears in three compounds rather than two and also because it appears as an element rather than a compound on the left side of the equation, which makes it easy to adjust the number of oxygen atoms without upsetting the balance of the other elements. Because there are ten O atoms on the product side of the equation, we need ten on the reactant side. We can easily achieve this by using 10/2 = 5 for the stoichiometric coefficient of the O2 reactant. This gives us the balanced chemical equation, 1 C3H8

+

5 O2

3 CO2

+

4 H2O

Normally, this would be written without showing the coefficient of one in front of the propane: C3H8 + 5 O2 : 3 CO2 + 4 H2O

Analyze Your Answer Our final equation has three carbon atoms, eight hydrogen atoms, and ten oxygen atoms on each side, so it is balanced.

Check Your Understanding The reaction we wrote above describes the complete combustion of propane. Under many conditions, though, hydrocarbons such as propane may not burn cleanly or completely, and other products may be formed. One such possible reaction is the formation of formaldehyde (CH2O) and water. Write a balanced chemical equation for a reaction in which propane and oxygen gas form CH2O and H2O as products.

MathConnections Usually, we balance chemical equations by the part systematic, part trial and error approach presented in Example 3.1. But it is fairly simple to transform that process into a more mathematical form, and doing so may help to emphasize the connection between balancing equations and the laws of conservation of matter.

70


Let’s illustrate the mathematical approach by reconsidering the equation we balanced in Example Problem 3.1. C3H8 +

O2 :

CO2 +

H2 O

It will help to assign variable names to each of the coefficients in the balanced equation. a C3H8 + b O2 : c CO2 + d H2O When the correct values of the coefficients (a, b, c, d ) are inserted, we know that the number of atoms of each element must be the same on both sides of the equation. We can express that idea as an equation for each element involved. Let’s start with carbon. On the left-hand side, we will have 3 carbon atoms for each molecule of C3H8, for a total of 3a carbon atoms. On the right side, we will have c carbon atoms. So we can set these two terms equal to one another. 3a = c Next we can write similar equations for the other elements. Looking at hydrogen, we get 8a = 2d. And for the oxygen, we have 2b = 2c + d. At this point, we have three equations for the four unknown coefficients, so we don’t appear to be in position to solve. The simplest way out of this predicament is just to set one coefficient arbitrarily equal to one. If we choose to set a = 1, we get the following: c = 3a = 3 2d = 8a = 8, and so d = 4 2b = 2c + d = 6 + 4 = 10, and so b = 5 Inserting these into the original equation gives us a balanced equation: C3H8 + 5 O2 : 3 CO2 + 4 H2O Because all of the resulting coefficient values are integers, we have found the smallest whole number coefficients that balance the equation, and our result is identical to that found in Example 3.1. If any fractional coefficients were obtained, we would simply multiply through by a constant to get whole numbers. This approach is quite general, and although it is not particularly popular with chemists, it has been used to write computer algorithms for balancing chemical equations.

For most people, the best way to learn how to balance chemical equations is just to practice until you gain confidence. The rules encountered for balancing equations are usually lists of steps that are prohibited rather than maps to be followed. For any reaction, there is more than one way to obtain the balanced equation, though some methods may be dramatically more difficult than others. A look at some common difficulties encountered by students new to balancing chemical equations may prove useful in developing successful strategies. One method that might appear to provide a route to the desired balanced chemical equation is to change the formulas of the molecules involved. Consider, for example, formaldehyde (CH2O) that might be formed during the burning or explosion of a hydrocarbon like propane. As long as there is still oxygen available, any formaldehyde


71

that is formed along the way can also burn. If CH2O burns incompletely, it will form carbon monoxide (CO) and water CH2O + O2 : CO + H2O Counting the number of atoms on each side of the equation reveals that we need one more oxygen on the product side. Many students find it tempting to write CH2O + O2 : CO2 + H2O This is undeniably a balanced chemical equation. In making this change, however, the entire meaning of the equation has been altered. Although this is a balanced equation, it does not depict the same reaction that we set out to describe. Carbon monoxide is no longer shown as a product, so this is not a valid way to balance the equation. Similarly, we can’t write the oxygen as O rather than O2, because the formaldehyde is reacting with oxygen molecules, not oxygen atoms. We may not change the chemicals involved in the equation, only their stoichiometric coefficients. So let’s take another look at our correct but unbalanced equation: CH2O + O2 : CO + H2O We still have two O atoms on the right and three on the left. One possible way to balance this equation is to multiply the O2 reactant by 1/2, so that there will be two oxygen atoms on each side: 1 CH2O + — O2 : CO + H2O 2 Generally, it is preferable to avoid using fractional coefficients because they could be interpreted as implying the existence of fractional molecules. But we can easily eliminate any fractions by multiplying the entire equation by a factor that removes them. In this case multiplying by 2 will do the trick, converting the 1/2 to 1. This brings us to a final form of the balanced equation: 2 CH2O

+

O2

2 CO

+

+

2 H2O

+

Note that we must multiply all of the stoichiometric coefficients by 2, not just the coefficient 1/2 in front of the oxygen. Otherwise, the balance that we have worked to establish would be destroyed. Some of these ideas are depicted pictorially in Figure 3.5.

3.3

Aqueous Solutions and Net Ionic Equations

For a chemical reaction to occur, the reactants involved must be able to come into contact with one another. In the gas or liquid phase, the molecules that make up the reactants move readily, allowing this contact to occur. In the solid phase, though, such motion is uncommon and reactions occur only very slowly if at all. One way to enable the needed contact between reactants that are solids under normal conditions is to dissolve them. Water is the most common medium for producing such solutions. Reactions that occur in water are said to take place in aqueous solution. To describe this important group of chemical reactions, first we need to define some key terms.

Solutions, Solvents, and Solutes Although water accounts for the vast majority of all the molecules in an aqueous solution, it is neither a reactant nor a product in many aqueous reactions. Water simply serves as the medium in which the reaction occurs. The entire liquid is called a solution, meaning

72


CH2O + O2

drawings here illustrate some common errors in balancing chemical equations.

Correct species, unbalanced equation

Change CO to CO2?

CH2O + O2

Figure 3.5 ❚ The molecular

CO + H2O

CO2 + H2O

Multiply O2 by 1/2 ?

CH2O + 1/2O2

CO + H2O

? Correct species, balanced equation, no valid molecular interpretation

Incorrect species, unbalanced equation Not valid

Undesirable Multiply ALL coefficients by 2?

Multiply O2 by 2? CH2O + O2

2 CH2O + O2

2 CO + 2 H2O

CO + H2O

Correct species unbalanced equation Not valid

Correct species, balanced equation, valid molecular interpretation Correct balanced form

that it is a homogeneous mixture of two or more substances. In any solution, the component present in greater amounts—usually much greater amounts—is called the solvent. In an aqueous solution, water is the solvent. Minor components of the solution are called solutes. The key feature of a solution is that the solutes dissolve in the solvent. When a substance dissolves, its particles are dispersed in the solvent, and typically many solvent molecules will surround any individual molecule or ion of solute. Although water is the most common solvent, it is not the only one. Not all solutions are even liquids. The same ideas can be applied to homogeneous mixtures of gases or solids. Air is a good example of a gaseous solution, and alloys like brass can be described as solid solutions. If we want to describe a pure substance, all we need to specify is its identity. If you pick up a bottle in the chemistry lab that is labeled “distilled water,” for example, you would know exactly what is in the container. But if the same container says only that it contains a solution of salt (NaCl) dissolved in water, you can’t be as sure what the contents are. Specifically, you would probably want to know something about the relative amounts of salt and water that were mixed to prepare the solution. So we need to specify at least one other key piece of information: the concentration. If there are many solute particles present, the solution is said to be concentrated. When there are few solute particles present the solution is called dilute. In Section 3.5, we will discuss various units that can be used to report solution concentrations.

3.3 Aqueous Solutions and Net Ionic Equations

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Photos: Lawrence S. Brown

Figure 3.6 ❚ In this sequence of photos, one of the authors prepares aqueous solutions of CuSO4. In the upper left panel, solid CuSO4!the solute!is transferred to a flask. In the upper right panel, water!the solvent!is added. The flask is shaken to speed the dissolution process (lower left). The final photo shows two CuSO4 solutions of different concentrations. The solution on the left has the higher concentration, as seen from its darker color.

In Chapter 12, we will see that solubility is not an “either or” process, but that there is continuous variability. These rules summarize the main quantitative tendencies we will learn then.

Some solutes can be dissolved in water to produce very concentrated solutions, whereas other substances do not dissolve to any measurable extent. So we can characterize various compounds in terms of their solubility in water. Those compounds that dissolve readily are said to be soluble, whereas those that do not dissolve significantly are called insoluble. Chemists often use a series of solubility rules to predict whether or not a particular compound is likely to dissolve in water. The most common rules are summarized in Table 3.1. Table

❚

3.1

Solubility guidelines for ionic compounds in water at room temperature

Usually Soluble

Exceptions

Group 1 cations (Li+, Na+, K+, Rb+, Cs+), ammonium (NH4+)

No common exceptions

Nitrates (NO3−), nitrites (NO2−)

Moderately soluble: AgNO2

Chlorides, bromides, iodides (Cl−, Br−, I−)

Insoluble: AgCl, Hg2Cl2, PbCl2, AgBr, Hg2Br2, PbBr2, AgI, Hg2I2, and PbI2

Fluorides (F−)

Insoluble: MgF2, CaF2, SrF2, BaF2, PbF2

Sulfates (SO42−)

Insoluble: BaSO4, PbSO4, HgSO4 Moderately soluble: CaSO4, SrSO4, Ag2SO4

Chlorates (ClO3−), perchlorates (ClO4−)

No common exceptions

Acetates (CH3COO−)

Moderately soluble: AgCH3COO

Usually Insoluble

Exceptions

Phosphates (PO43−)

Soluble: (NH4)3PO4, Na3PO4, K3PO4

2−

74

Carbonates (CO3 )

Soluble: (NH4)2CO3, Na2CO3, K2CO3

Hydroxides (OH−)

Soluble: LiOH, NaOH, KOH, Ba(OH)2 Moderately soluble: Ca(OH)2, Sr(OH)2

Sulfides (S2−)

Soluble: (NH4)2S, Na2S, K2S, MgS, CaS



3.2

Which of the following compounds would you predict are soluble in water at room temperature? (a) KClO3, (b) CaCO3, (c) BaSO4, (d) KMnO4

Strategy Solubility guidelines for common ions are given in Table 3.1. So we will identify the ions in each compound and consult the table as needed to determine the solubilities.

Solution (a) KClO3 is potassium chlorate. From the solubility guidelines in Table 3.1, we see that compounds containing K+ and ClO3− tend to be soluble and that no common exceptions are mentioned. So we predict that KClO3 should be soluble. (b) CaCO3 is calcium carbonate. Again consulting the table, we see that carbonates are generally insoluble and that CaCO3 is not listed among the exceptions. Thus CaCO3 should be insoluble. (c) BaSO4 is barium sulfate. Although most sulfates are soluble, BaSO4 is listed in the table as an exception to that rule. Therefore we expect BaSO4 to be insoluble. (d) KMnO4 is potassium permanganate. Although the permanganate ion (MnO4−) is not listed in Table 3.1, the table does tell us that all compounds of K+ are soluble. So we would predict that KMnO4 should be soluble.

Discussion Here we simply consulted the table to check the solubility of each compound. Chemists generally gain familiarity with many of these solubility rules and learn to recognize soluble and insoluble salts without consulting such a table. You should check with your instructor to see whether you are expected to memorize these rules and exceptions.

Check Your Understanding Which of the following compounds would be soluble in water at room temperature? (a) NH4Cl, (b) KOH, (c) Ca(CH3COO)2, (d) Ba3(PO4)2

Although these rules imply that solubility is a simple yes or no question, the reality is more complicated than that. Saying that a compound is soluble does not mean that we could dissolve limitless amounts of it in a small beaker of water. If we keep adding more of the solute, eventually we will observe that the added material does not dissolve. When this occurs we have established a saturated solution. The concentration at which a given solution will become saturated depends on the identity of the solute, so it is useful to have a quantitative measure of solubility. The units used to report this number can vary, but one common choice is mass of solute per 100 g of solvent. For example, the solubility of table salt in water at room temperature is 35.7 g NaCl/100 g H2O. This corresponds to about one-third of a cup of salt in a cup of water. Notice that solubility gives us a ratio, which we could use in the same way as we did the mass density in Example Problem 1.6 on page 20. One final observation to make about solutions is that many ionic compounds will dissociate into individual ions when they dissolve in water. Thus the salt solution above actually contains Na+ and Cl− ions, not NaCl molecules. The availability of freely moving charges allows these solutions to conduct electricity. Any substance that dissolves in water to produce an aqueous solution that conducts electricity is called an electrolyte. Substances whose solutions do not conduct electricity are called nonelectrolytes. We can divide electrolytes further into two groups. Strong electrolytes dissociate completely, so that only individual ions are present in the solution, with virtually no intact molecules. In contrast, weak electrolytes dissociate only


75

Photos: © Cengage Learning/Charles D. Winters

H2O molecule

(c)

(b)

(a)

Sugar molecule

Hydrogen ion, H+

Acetic acid molecule, CH3COOH

Acetate ion, CH3COO –

Potassium ion, K+

Chromate ion, CrO42–

Figure 3.7 ❚ The photos show a classroom demonstration in which a pair of copper rods is dipped into different solutions. If the solution conducts electricity, a circuit is closed and the lightbulb lights. (a) Ordinary sugar (C12H22O11) is a nonelectrolyte, so a solution of sugar dissolved in water does not conduct electricity. (b) Acetic acid (CH3COOH) is a weak electrolyte. An acetic acid solution contains low concentrations of ions and conducts electricity well enough to light the bulb dimly. (c) Potassium chromate (K2CrO4) is a strong electrolyte. A potassium chromate solution contains higher concentrations of ions and conducts electricity well enough to light the bulb brightly. The illustrations accompanying each photo emphasize the solute species in each solution.

partially; their solutions contain both intact molecules and individual ions in measurable quantities. Figure 3.7 shows the differences between these classes of solutes.

Chemical Equations for Aqueous Reactions We can describe the process by which a compound dissolves in water by a chemical equation. When a covalently bonded material such as sugar dissolves in water, the molecules remain intact: C6H12O6(s) : C6H12O6(aq) By contrast, when an ionic solid dissolves in water, it breaks into its constituent ions. This is called a dissociation reaction. The dissolving of sodium chloride discussed above is a common example of this process: NaCl(s) : Na+(aq) + Cl−(aq) In both of these chemical equations, the water molecules are not shown explicitly, although their presence is indicated by the “(aq)” on the product side. This omission reflects the general tendency of a solvent to dissolve a solute but not to react chemically with it. When water appears in a chemical equation, the reaction involves water as either a reactant or product. In addition to these fairly simple reactions by which we describe compounds dissolving in water, many important reactions take place in water. The chemical equations we write to describe these reactions can be written in any of three forms; the choice of equation is based mainly on the context in which the equation is used. We’ll

76


illustrate these three types of equations for a reaction that is important in the production of commercial explosives: the synthesis of ammonium nitrate from ammonia and nitric acid. Ammonium nitrate is a precursor to many important explosives and also has considerable explosive potential itself. Because ammonium nitrate is also widely used as a fertilizer, it is readily available for purchase. An unfortunate result of this availability is that it has often been used in bombings, including the 1995 attack on the Murrah Federal Building in Oklahoma City. In the commercial preparation of ammonium nitrate, pure ammonia (NH3) in the gas phase is combined with concentrated aqueous nitric acid (HNO3). This produces a highly concentrated aqueous solution of ammonium nitrate, which is then dried and formed into small pellets (called prills). The molecular equation shows the complete formula of each compound involved. Both intact and dissociated compounds are shown, followed by the (aq) to designate an aqueous solution when appropriate: HNO3(aq) + NH3(g) : NH4NO3(aq) Frequently, the compounds involved in aqueous chemistry are ionic, as is true for both HNO3 and NH4NO3 in this example. These ionic compounds dissociate when they dissolve in water, so the actual solution does not contain intact molecules of these compounds. This dissociation is emphasized if we write the total ionic equation rather than the molecular equation. This form emphasizes what is actually present in the reacting mixture by writing dissociated compounds as separated ions in the solution. The total ionic equation for our example reaction from above is thus H+(aq) + NO3−(aq) + NH3(g) : NH4+(aq) + NO3−(aq) The H+ ion is rather special because it is present in all aqueous solutions of acids. Strictly speaking, an H+ ion would be nothing more than a proton. But as we will see in more detail later, such a bare proton in aqueous solution would actually be tightly surrounded by water molecules. Thus the H+ ion is frequently written as H3O+ to remind us of this fact. If we write the ion in this way, our total ionic equation becomes H3O+(aq) + NO3−(aq) + NH3(g) : NH4+(aq) + NO3−(aq) + H2O(,) Note that now we must include a molecule of water on the right-hand side to balance the atoms that we added when we switched our representation of the H+ ion from H+ to H3O+. Careful inspection of this equation shows that one ion, NO3−(aq), appears on both sides. Thus it is not an active participant in the reaction. Ions that are uninvolved in the chemistry are referred to as spectator ions. In many instances, we are not particularly concerned with these spectator ions and need not include them in the chemical equation. When the spectator ions are left out, the result is a net ionic equation. Using the same reaction as an example, we would have either

The process of removing the spectator ions from the equation is similar to what you would do if you encountered an algebraic equation that had the same term on both sides.

H+(aq) + NH3(g) : NH4+(aq) or H3O+(aq) + NH3(g) : NH4+(aq) + H2O(,) This depends on whether we choose to write H+ or H3O+ to represent the cation from the nitric acid. Because we have removed anions from both sides of the equation, we now have a net charge of 1+ on each side. This is one way to recognize that you are looking at a net ionic equation: there is often a nonzero net charge on each side of the equation. If the equation is properly balanced, though, the net charge on each side must be the same. This rule reflects the fact that electrical charge must be conserved. At this point, you are probably wondering, “Which of these equations is the correct one?” The answer is that all of these forms provide valid descriptions of the reaction, so none is intrinsically “better” than the others. In this particular case, we would probably not choose the net ionic equation because the identity of the anion (NO3−)


77

is important if our aim is to produce solid ammonium nitrate for the explosives industry. But in many other cases, we might have little or no interest in the identity of a spectator ion, so the simplicity of the net ionic equation might make it attractive. The total ionic equation is more cumbersome and is used less often.

Acid–Base Reactions Two especially important categories of aqueous solutions are acids and bases. Examples of them are easy to find in our everyday lives as well as in the chemical industry. For our present purposes, we will define an acid as any substance that dissolves in water to produce H+ (or H3O+) ions and a base as any substance that dissolves in water to produce OH− ions. Table 3.2 lists some common acids and bases. Like other solutes, acids and bases can be either strong or weak electrolytes. Strong acids or bases dissociate completely in water, so that the resulting solution contains essentially no intact solute molecules. We can write the following chemical equations for the dissociation of HCl and NaOH: Unlike H+ ions, hydroxide ions are not represented as having combined with water from the solvent, so water doesn’t appear in the equation for a strong base.

We will revisit the concepts of acids and bases in Chapter 12 and provide additional detail at that time.

Strong acid: HCl(g) + H2O(ℓ) : H3O+(aq) + Cl−(aq) Strong base: NaOH(s) : Na+(aq) + OH−(aq) The only common strong acids and bases are those shown in Table 3.2. Weak acids or bases undergo only partial ionization, so that their solutions contain intact molecules as well as dissociated ions. When we write equations for the dissolution of weak electrolytes like these, we use a “two-way” arrow that emphasizes that the reaction does not proceed completely from left to right. For common weak acids, it is relatively easy to write the needed ionization equation. Many weak acids contain the !COOH functional group, and the H atom from that group tends to be lost in solution. Acetic acid is a good example. Weak acid: CH3COOH(aq) + H2O(ℓ) N H3O+(aq) + CH3COO−(aq) Table

❚

3.2

Strong and weak acids and bases

Strong Acids

Strong Bases

HCl

Hydrochloric acid

LiOH

Lithium hydroxide

HNO3

Nitric acid

NaOH

Sodium hydroxide

H2SO4

Sulfuric acid

KOH

Potassium hydroxide

HClO4

Perchloric acid

Ca(OH)2

Calcium hydroxide

HBr

Hydrobromic acid

Ba(OH)2

Barium hydroxide

HI

Hydriodic acid

Sr(OH)2

Strontium hydroxide

Weak Acids

Weak Bases

H3PO4

Phosphoric acid

NH3

Ammonia

HF

Hydrofluoric acid

CH3NH2

Methylamine

CH3COOH

Acetic acid

HCN

Hydrocyanic acid

Note: All common strong acids and bases are shown, but only representative examples of weak acids and bases are listed.

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For weak bases, the situation is slightly less obvious. Unlike the strong bases listed in Table 3.2, most weak bases do not contain !OH groups. So it may not be obvious that dissolving them will produce hydroxide ions in solution (i.e., that they are bases at all). The most common weak base is ammonia (NH3), which reacts in water according to the following equation: Weak base: NH3(aq) + H2O(ℓ) N NH4+(aq) + OH−(aq) Again we use the two-way arrow to show that intact NH3 molecules as well as NH4+ and OH− ions will be present in the solution. Many other weak bases are amines. They can be thought of as derivatives of ammonia in which one or more of the H atoms have been replaced by methyl groups or longer hydrocarbon chains. Because many acidic and basic solutions occur in nature, observations about acids and bases date back hundreds of years. One of the most important observations is that a solution cannot be both acidic and basic at the same time. Mixing an acid and a base leads to a reaction known as neutralization, in which the resulting solution is neither acidic nor basic. To understand the origins of this neutralization, we might start with the definitions of acid and base. Acid solutions contain H3O+ ions, and bases contain OH− ions. So for a solution to be both an acid and a base simultaneously, it would need to contain both of these species. Looking at those ions, it should be easy to see why this is not feasible. Hydronium ions and hydroxide ions combine readily to form water: H3O+(aq) + OH−(aq) : 2 H2O(,) This reaction will take place whenever an acid and a base are combined and will always prevent a solution from being both acidic and basic simultaneously.


3.3

When aqueous solutions of acetic acid and potassium hydroxide are combined, a neutralization reaction will occur. Write molecular, total ionic, and net ionic equations for this process.

Strategy This is the reaction of a weak acid with a strong base. The reaction between an acid and a base always produces water as a product, along with an ionic compound formed from the remaining ions. (This second product is often called a salt.) We can begin by using this idea to write the molecular equation. To generate the ionic equations, then, we account for the dissociation of strong electrolytes into their constituent ions.

Solution A hydrogen atom from the acid and the hydroxide ion from the base will

produce water. The remaining ions will be an acetate anion (CH3COO− from the acid) and a potassium ion (K+, from the base). Combining these gives us potassium acetate, KCH3COO. This lets us write the molecular equation: CH3COOH + KOH : H2O + KCH3COO To produce the total ionic equation, we must determine which of these species to write as dissociated ions. According to Table 3.2, acetic acid is a weak acid. This means that it is only partially dissociated in solution, and so we write it as an intact molecule. KOH, on the other hand, is a strong base, and so it will dissociate completely. So we write this as a pair of ions, K+ and OH−. According to Table 3.1, both potassium ions and acetate ions tend to produce soluble compounds. So, KCH3COO will be soluble and hence will dissociate into its constituent ions. Putting all of this together lets us write the total ionic equation: CH3COOH(aq) + K+(aq) + OH−(aq) : H2O(ℓ) + K+(aq) + CH3COO−(aq)


79

Looking at this equation, we see potassium ions on both sides. Thus potassium is a spectator ion and can be deleted to give the net ionic equation: CH3COOH(aq) + OH−(aq) : H2O(ℓ) + CH3COO−(aq)

Check Your Understanding Write molecular, total ionic, and net ionic equations for the reaction between hydrochloric acid and calcium hydroxide.

A number of commercial products are available for cleaning up acid spills, and they all rely on neutralization. Typically a mixture of compounds such as calcium carbonate, magnesium oxide, and sodium carbonate is employed, often together with a dye that changes color when the acid spill has been neutralized. Neutralization is only one example of aqueous reaction chemistry. Another common type of reaction is the precipitation reaction, in which a solid (called the precipitate) is formed and drops out of solution as shown in Figure 3.8. Example Problem 3.4 provides an example of the three forms of chemical equations for precipitation reactions.

E XAM PL E PRO BLEM

3 .4

When aqueous sodium carbonate is added to barium chloride, the solution turns cloudy white with solid barium carbonate. Write the molecular, ionic, and net ionic equations for this reaction.

Strategy First we will use the nomenclature rules that we presented in Chapter 2 to write formulas for all of the compounds involved. Then we construct a balanced chemical equation using complete formula units. Next we must identify any ionic compounds that are present, and we write them as dissociated ions to give the total ionic equation. Finally, to obtain the net ionic equation, we must eliminate any spectator ions from both sides of the equation.

Solution Formulas for the substances listed are Na 2CO3, BaCl2, and BaCO3. Although BaCO3 is the only product listed in the problem statement, the sodium and chlorine atoms must also appear somewhere on the right-hand side of our equation. So NaCl will also be produced as a product. We write reactants on the left and products on the right to set up the molecular equation and then balance: Na2CO3(aq) + BaCl2(aq) : BaCO3(s) + 2 NaCl(aq)

Photos: Lawrence S. Brown

Figure 3.8 ❚ The photos show a demonstration in which clear and colorless solutions of KI and Pb(NO3)2 are mixed and react to form a precipitate of PbI2. (The precipitate is bright yellow.) The accompanying illustration shows a molecular scale representation of the individual reactant solutions and the final product.

I–

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K+

NO3–

Pb2+

Next we recognize that the three substances denoted as “(aq)” in that equation will exist as dissociated ions rather than intact molecules. This leads us to the total ionic equation: 2 Na+(aq) + CO32−(aq) + Ba2+(aq) + 2 Cl−(aq) : BaCO3(s) + 2 Na+(aq) + 2 Cl−(aq) Finally, we can cancel the sodium and chloride ions because they appear in equal numbers on both sides of the equation. This leads us to the net ionic equation: CO32−(aq) + Ba2+(aq) : BaCO3(s)

Check Your Understanding When sodium sulfate, Na2SO4, reacts with lead nitrate, Pb(NO3)2, the products are solid lead sulfate and aqueous sodium nitrate. Write the molecular, ionic, and net ionic equations for this reaction. Note that for a precipitation reaction, the net ionic equation does not have a net charge on either the reactant or product side.

3.4


The chemical equations that we have been learning to write are symbolic descriptions of chemical reactions. But to interpret these equations, we must think about them from another point of view, in terms of the actual substances and processes they represent. As often happens in chemistry, we can do this at either the microscopic or the macroscopic level. The microscopic interpretation visualizes reactions between individual molecules, and that interpretation is the one we have used thus far. The macroscopic interpretation pictures reactions between bulk quantities. Neither perspective is inherently better than the other; we simply use the point of view that is best suited to whatever circumstances we are considering. To make the connection between the two worlds, we need to look at the way chemists count the molecules in a large sample.

Interpreting Chemical Equations In describing how to write and balance a chemical equation, we adopted a way of looking at reactions that seems natural. We spoke of one molecule reacting with another molecule to form some new compound that we called the product. Real chemical reactions rarely involve just one or two molecules though, so we will need to think of chemical equations in terms of larger quantities, too. Let’s return to our first example, the explosive reaction between hydrogen and oxygen to produce water: 2 H2(g) + O2(g) : 2 H2O(g) We have read this chemical equation as saying that “two H2 molecules react with one O2 molecule to form two H2O molecules.” But suppose we carried out the same reaction starting with 20 H2 molecules rather than two. The balanced equation says that we need one oxygen molecule for every two H2 molecules. So in this case, we could say that 20 H2 molecules react with 10 O2 molecules to form 20 H2O molecules. Two key features of these statements must be emphasized. First, both statements are correct because they use the ratios established by the stoichiometric coefficients. There are twice as many hydrogen molecules as there are oxygen molecules in each case. Second, both statements refer to numbers of particles. When we interpret a chemical reaction, we must remember that the stoichiometric coefficients always refer to numbers of particles. In the example we just considered, the numbers were 1 and 2 or 10 and 20. We could just as easily have said 1,000,000 and 2,000,000 or 45,600,000 and 91,200,000. Provided we use numbers that satisfy the correct stoichiometric ratio, we will always have a reasonable description of an actual chemical reaction.

3.4


81

Avogadro’s Number and the Mole

Because of its fundamental importance, Avogadro’s number has been measured extremely carefully. The internationally accepted value is actually 6.02214179 × 10 23.

When we begin to consider chemical reactions with macroscopic amounts of materials, the numbers of molecules involved become staggeringly large. But people generally don’t like to work with huge numbers, because it is very difficult to develop an intuitive feeling for them. So inevitably, when confronted with the fact that any real sample contains an almost preposterous number of molecules, scientists invented a way to make the numbers manageable. The simple way to do this is just to count the molecules in terms of some large quantity rather than individually. The particular quantity that chemists have chosen to count by is called the mole. One mole is defined as the number of atoms in exactly 12 grams of 12C. This number is also referred to as Avogadro’s number, and its value is 6.022 × 1023. When chemists want to count molecules, they count by moles, so that the resulting numbers will be easier to work with. There are several reasons why this number is useful in chemistry. First, it is a large number that gives a useful unit for counting particles as small as atoms. In this sense, the mole is to a chemist what a dozen is to a baker. Second, one mole is the same number of particles no matter what substance we are talking about. So a mole of H2 contains the same number of particles as a mole of TNT (trinitrotoluene, C7H5N3O6); each has 6.022 × 1023 molecules. You may be thinking that the value of Avogadro’s number seems like a rather peculiar choice for a counting unit. Why not just make it an even 1023, for example, and simplify some calculations? The answer is that the chosen value of Avogadro’s number produces a very convenient connection between the mass of an atom or molecule and the mass of a mole of those atoms or molecules. Our definition of the mole said that a mole of 12C has a mass of exactly 12 g. Earlier, we said that an atom of 12C has a mass of exactly 12 amu. This is not merely a coincidence, but rather a very deliberate decision. Once we have defined the amu and the gram independently, Avogadro’s number is fixed as the number of amu in one gram. It may help to envision a hypothetical experiment: suppose that we had a balance that could accurately measure the mass of individual atoms and a way of loading atoms onto that balance one at a time. We could then determine Avogadro’s number by loading carbon-12 atoms onto the balance, one by one, until the mass read exactly 12 grams. Of course, no such experiment is possible, so Avogadro’s number has been measured by more indirect methods. The mass of 6.022 × 1023 atoms of any element is the molar mass of that element, and its value in grams per mole is given in most periodic tables. The molar mass of silicon, for example, is greater than that of carbon because an atom of silicon is heavier than an atom of carbon. And the molar mass of each element takes into account the natural isotopic abundances. So the molar mass of carbon is 12.011 g/mol, reflecting the weighted average between the masses of 12C and 13C in the same way as we discussed in Chapter 2.

Charles Steele

Figure 3.9 ❚ One mole samples of various elements are shown. Back row (left to right): bromine, aluminum, mercury, and copper. Front row (left to right): sulfur, zinc, and iron.

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The mole is the key to the macroscopic interpretation of a chemical equation. Consider the same equation we have been discussing: 2 H2(g) + O2(g) : 2 H2O(g) If we want to read this in terms of moles, we can say “two moles of H2 and one mole of O2 react to form two moles of H2O.” Because each mole contains the same number of molecules, the 2:1 mole ratio between the reactants is the same as the 2:1 ratio for the numbers of molecules. Chemical equations and their stoichiometric coefficients always provide ratios of numbers of particles, not masses.

Determining Molar Mass Because balanced chemical equations are always expressed in terms of numbers of particles, it will be convenient to have some easy way to determine the number of particles in a given sample of some substance. But there is no simple laboratory instrument that can measure the number of moles in a sample directly. Instead, we usually determine the number of moles indirectly from the mass of a sample. The link between the mass of a sample and the number of moles present is the molar mass of the substance in question. To determine the molar mass of a compound, we can exploit the idea of conservation of mass. Consider one mole of water as an example. We know that one mole of the compound must contain Avogadro’s number of H2O molecules. Furthermore, we know that each of those molecules must contain one O atom and two H atoms. Avogadro’s number of oxygen atoms is one mole, and we know from the periodic table that one mole of O atoms has a mass of 16.0 g. Because each molecule contains two hydrogen atoms, the entire one mole sample will contain two moles of H atoms. Again consulting the periodic table, we see that a mole of H atoms has a mass of 1.0 g, so two moles must have a mass of 2.0 g. The masses of the O and H atoms must be the same as the mass of the mole of H2O, so we can simply add them to get 18.0 g as the mass of one mole of H2O. In other words, the molar mass of H2O is 18.0 g/mol. We can do the same thing for any compound; the sum of the molar masses of all of the atoms is the molar mass of the compound. In Example Problem 3.5, we determine the molar mass of several explosive compounds.


For simplicity, we use whole numbers for the molar masses of hydrogen and oxygen here. In general, you should use molar masses with at least as many significant figures as the actual data involved in the problem you are working.

3.5

Determine the molar mass of each of the following compounds, all of which have been used as explosives: (a) lead azide, PbN6, (b) nitroglycerin, C3H5N3O9, (c) mercury fulminate, Hg(ONC)2

Strategy We must determine the mass contributed by each element and then add them up to calculate the molar mass. When parentheses appear in the formula, each atom inside the parentheses must be multiplied by its own subscript and by the subscript appearing after the right-hand parenthesis.

Solution (a) PbN6:

1 mole Pb: 6 moles N:

1 × 207.2 g/mol = 207.2 6 × 14.0067 g/mol = 84.0402 Molar mass = 291.2

g/mol g/mol g/mol

(b) C3H5N3O9:

3 moles C: 5 moles H: 3 moles N: 9 moles O:

3× 5× 3× 9×

g/mol g/mol g/mol g/mol g/mol

12.011 g/mol 1.0079 g/mol 14.0067 g/mol 15.9994 g/mol Molar mass

= 36.033 = 5.0395 = 42.0202 = 143.995 = 227.087

3.4


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(c) Hg(ONC)2:

1 mole Hg: 2 moles O: 2 moles N: 2 moles C:

1 × 200.59 g/mol 2 × 15.9994 g/mol 2 × 14.0067 g/mol 2 × 12.011 g/mol Molar mass

= 200.59 = 31.9988 = 28.0134 = 24.022 284.62

g/mol g/mol g/mol g/mol g/mol

Analyze Your Answer Looking at these three compounds, we see that nitroglycerin has the smallest molar mass even though it clearly contains the most atoms per molecule. Does this make sense? The key factor is that both lead azide and mercury fulminate contain elements with large molar masses, so it is not surprising that nitroglycerin is the least massive of these three molecules.

Discussion The number of significant figures to which the molar masses of the elements are known varies from one element to another, as you can see if you examine the periodic table inside the back cover of this book. So the appropriate number of significant figures in the molar masses here varies according to the elements in each molecule.

Check Your Understanding What are the molar masses of the following compounds, which are used in the preparation of various explosives? (a) H2SO4, (b) HNO3, (c) (NH2)2HNO3

3.5


We’ve seen the importance of the concept of the mole when dealing with chemical reactions involving macroscopic quantities of material. But generally it is not possible to measure the number of moles in a sample directly because that would imply that we can count molecules. Molar masses provide a crucial connection and allow us to convert from masses, which can be measured easily, to numbers of moles. The mass and the number of moles are really just two different ways of expressing the same information—the amount of a substance present. The molar mass functions much like a unit conversion between them, allowing us to go from mass to moles or from moles to mass. The following examples illustrate typical uses of molar mass.

E XAM PL E PRO BLEM

3 .6

A sample of the explosive TNT (C7H5N3O6) has a mass of 650.5 g. How many moles of TNT are in this sample? How many molecules is this?

Strategy We are asked to convert from mass to moles. So we must determine the molar mass of the substance and then use it to carry out the conversion. Once we know the number of moles, we can easily find the number of molecules because we know that one mole contains Avogadro’s number of molecules. Trinitrotoluene (TNT), C7H5N3O6

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Solution First we will calculate the molar mass of TNT. C7H5N3O6:



7 × 12.011 g/mol 5 × 1.0079 g/mol 3 × 14.0067 g/mol 6 × 15.9994 g/mol Molar mass

= = = = =

84.077 g/mol 5.0395 g/mol 42.0201 g/mol 95.9964 g/mol 227.133 g/mol

Now we can use the molar mass to convert from mass to moles: 1 mol TNT 650.5 g TNT × ——————— = 2.864 mol TNT 227.133 g TNT Finally, we can convert from moles to molecules using Avogadro’s number: 6.022 × 1023 molecules TNT 2.864 mol TNT × ————————————— = 1.725 × 1024 molecules TNT 1 mol TNT

Analyze Your Answer Admittedly, it is difficult to have an intuitive feeling for the number of molecules in any sample. In assessing our answer, we can begin by noting that it is a very large number. This makes sense because we are calculating the number of molecules in a macroscopic amount of TNT. We could also check a little more closely by looking at the intermediate step in which we found the number of moles. Because the molar mass of TNT is slightly larger than 200 g/mol, we should have roughly 3 moles in the 650-g sample. This is consistent with the value we found.

A common student error in this type of calculation is to divide by Avogadro’s number rather than multiply. This leads to a very small numerical answer. A number of molecules less than one is not possible, of course, and is a sure sign of an error in the calculation.

Discussion The mass of this sample corresponds to a little less than 1.5 pounds, so the sample would probably be roughly the size of a brick. Note the enormous number of molecules in this fairly small amount of material.

Check Your Understanding Liquid nitromethane was once widely used as an explosive but is now much more common as a fuel for high performance engines in drag racing. How many moles are present in 1.00 gallon of pure nitromethane (CH3NO2), which has a density of 1.137 g/mL? In some cases, we might calculate the number of moles of a substance needed in a particular reaction. Then, if we want to carry out that reaction, we would need to know what mass of sample to prepare. We can easily invert the process above to convert from the number of moles to mass.


3.7

A demolition engineer is planning to use the explosive ethylenedinitramine (C2H6N4O4, also known as halite) to bring down an abandoned building. Calculations show that 315 moles of the compound will provide the necessary explosive force. How many pounds of C2H6N4O4 should be used?

Strategy We are asked to convert from a number of moles to a mass. We know that the link between these two quantities is the molar mass of the compound, so we can start by calculating that. Once we know the molar mass, we can use it to find the required mass of explosive. Finally, we can convert that result from grams to pounds.

Solution We will need to calculate the molar mass of the compound. 2 × 12.011 g/mol = 24.022 g/mol 6 × 1.0079 g/mol = 6.0474 g/mol 4 × 14.0067 g/mol = 56.0268 g/mol 4 × 15.9994 g/mol = 63.9976 g/mol Molar mass = 150.094 g/mol Now we can use the molar mass and the required number of moles to find the appropriate mass. C2H6N4O4:


150.094 g C2H6N4O4 315 moles C2H6N4O4 × —————————— = 4.73 × 104 g C2H6N4O4 1 mole C2H6N4O4


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Finally, we can convert this from grams to pounds: 1 pound 4.73 × 104 g C2H6N4O4 × ———— = 104 pounds C2H6N4O4 needed 453.59 g

Analyze Your Answer We can assess the reasonableness of this answer by comparing it to Example Problem 3.6. For these two problems, the molar masses of the compounds are roughly comparable. In this problem, we have about 300 moles, whereas in the previous one, we had about 3 moles, so it’s not surprising that our answer is close to 100 times larger than that in our last exercise.

Check Your Understanding Approximately 80 billion moles of ammonium nitrate (NH4NO3) were produced globally in 2008. How many tons of ammonium nitrate is this?

Elemental Analysis: Determining Empirical and Molecular Formulas When a new molecule is synthesized, an elemental analysis is routinely performed to help verify its identity. This test, which measures the mass percentage of each element in the compound, is also frequently done as part of the process of identifying any substance whose composition is unknown. The mass percentages describe the compound’s composition, and so they must be related to its chemical formula. But the data obtained from elemental analysis describe the composition in terms of the mass of each element, whereas the formula describes the composition in terms of the number of atoms of each element. So these are two different representations of very similar information, and the molar masses of the elements provide a connection between them. The process of obtaining the empirical formula of a compound from its percent composition by mass is best illustrated by an example.

E XAM PL E PRO BLEM

3 .8

The explosive known as RDX contains 16.22% carbon, 2.72% hydrogen, 37.84% nitrogen, and 43.22% oxygen by mass. Determine the empirical formula of the compound.

Strategy The empirical formula is based on the mole ratios among the elements in the compound, and the data given are in terms of mass. Molar mass provides the link between mass and moles, as usual. We can begin by choosing a convenient mass of the compound—usually 100 g—and use the percentages given to find the mass of each element in that sample. Then we will convert those masses into numbers of moles of each element. The ratios between those numbers of moles must be the same for any sample of the compound. Finally, we will need to convert those ratios into whole numbers to write the empirical formula.

Solution Consider a 100-g sample of RDX. From the percentages given, that sample will contain 16.22g C, 2.72g H, 37.84g N, and 43.22g O. We will convert each of these masses into moles: 1 mol C 16.22g C × ————— = 1.350 mol C in 100 g RDX 12.011 g C 1 mol H 2.72g H × ————— = 2.70 mol H in 100 g RDX 1.0079 g H

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1 mol N 37.84g N × —————— = 2.702 mol N in 100 g RDX 14.0067 g N 1 mol O 43.22g O × —————— = 2.701 mol O in 100 g RDX 15.9994 g O These numbers tell us the ratios of C:H:N:O in the compound. We might be tempted to write a formula of C1.35H2.70N2.70O2.70. But we know that the correct empirical formula requires whole numbers of each element. In this case, it may be easy to see the correct ratios, but we need to approach such a problem systematically. The usual strategy is to divide all of the numbers of moles we obtained above by the smallest number of moles. This assures that the smallest resulting number will always be one. In this case, we will divide all four numbers of moles by 1.350 because it is the smallest of the four. 1.350 mol C 1.350

—————— = 1

2.70 mol H 1.350

—————— = 2

2.702 mol N 1.350

—————— = 2.001 ≈ 2

2.701 mol O 1.350

—————— = 2.001 ≈ 2

The result is a small whole number ratio: 1 mole C:2 moles of H:2 moles N:2 moles O. So the empirical formula is CH2N2O2. The mass percentage data allow us to determine only the empirical formula, not the molecular formula, because we do not have any information about how large or small a molecule of the compound might be.

Analyze Your Answer We know that the coefficients in a chemical formula must be whole numbers. The fact that all four of the coefficients we found are so close to integers means that our proposed formula is plausible.

Discussion In working with this type of problem, it is a good idea to use accurate molar masses for the elements and to carry as many significant figures as possible throughout the calculation. This ensures that no large rounding errors occur and helps make it easy to decide whether the final coefficients obtained are whole numbers. The steps we used above will produce whole numbers for all of the elements only if at least one of the subscripts in the empirical formula is a one. Otherwise, the result for one or more elements will still not be an integer. In nearly all such cases, though, such nonintegral results will be readily identified as small rational fractions (e.g., 1.5 or 2.33). We can obtain integral values by multiplying all of the coefficients by an appropriate integer. This is demonstrated in the “Check Your Understanding” exercise that follows.

Check Your Understanding Nitroglycerin contains 15.87% C, 2.22% H, 18.50% N, and 63.41% O by mass. Determine the empirical formula of the compound.

If we also know the molar mass of the compound, the process used in the previous example can easily be extended to allow the determination of the molecular formula. The connection between mass and moles can also be applied in other ways. Materials scientists and engineers often describe the composition of alloys in terms of either weight percentage (wt %) or mole percentage (mol %). Converting between these two units involves the use of molar mass, as shown in Example 3.9.

An alloy is a solution in which both the solute and solvent are solids. Materials engineers can choose or design alloys to have useful properties for a particular application.


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E XAM PL E PRO BLEM

3 .9

Alloys of palladium and nickel are often used in the manufacture of electronic connectors. One such alloy contains 70.8 mol % Pd and 29.2 mol % Ni. Express the composition of this alloy in weight percentage.

Strategy In order to get started, we need to choose an amount of the alloy. Although any amount would work, one mole might be convenient. Once we have specified a number of moles, we can use molar masses to determine the mass of each metal and the total mass. And from those we can calculate the weight percentages.

Solution The molar masses of palladium and nickel are 106.42 g/mol and 58.69 g/mol, respectively. From the given mole percentages, we know that one mole of alloy would contain 0.708 mole of Pd and 0.292 mole of Ni. We can use these values to determine the mass of each component in one mole of alloy. 106.42 g Pd mPd = 0.708 mol Pd × —————— = 75.4 g Pd 1 mol Pd 58.6934 g Ni mNi = 0.292 mol Ni × —————— = 17.1 g Ni 1 mol Ni So the total mass of one mole of alloy is just 75.4 g Pd + 17.1 g Ni = 92.5 g alloy Finally, calculate the percentage by mass (weight) of each metal. 75.4 g Pd wt % (Pd) = —————— × 100% = 81.5 wt % Pd 92.5 g alloy 17.1 g Ni wt % (Ni) = —————— × 100% = 18.5 wt % Ni 92.5 g alloy

Analyze Your Answer It makes sense that the weight percentage of the palladium is higher than the mole percentage because palladium has a higher molar mass than nickel.

Discussion There is nothing special about the assumption of one mole in this problem. It was chosen because it was likely to provide numbers for the masses that were a “comfortable” size to think about. If the problem were worked in the other direction (from wt % to mol %), we might choose 100 g as the amount of material to use in the calculation. An additional unit sometimes used for alloy composition is atomic percentage (at %). The conversion between at % and wt % involves the same steps as the mol % and wt % conversion in this problem.

Check Your Understanding Eighteen carat gold typically contains 75% gold, 16% silver, and 9% copper by weight. Express this composition in mol percentages.

Molarity Molar mass is useful for chemical calculations because it provides a connection between a quantity that is easy to measure (mass) and one that is conceptually important (moles). Another readily measured quantity is volume. When we work with aqueous solutions, we often use volume in calculations rather than mass. It should not be surprising, therefore, that we would want to define quantities that will help us relate a volume measurement to the number of moles. Many different ways to express the concentration of a solute in a solvent have been developed. To define any unit of concentration, we must know both the amount of solute and the quantity of solvent in the solution. The most commonly used concentration unit

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in chemistry is molarity or molar concentration. Molarity, represented by the symbol M, is defined as the number of moles of solute per liter of solution: moles of solute Molarity (M) = ——————— liter of solution

(3.1)

The definition of molarity points to the ways in which it can be used. It provides a relationship among three things: molar concentration, moles of solute, and liters of solution. If we know any two of these quantities, we can determine the third one. We can measure the volume of a solution in the laboratory. Multiplying that volume (expressed in liters) by molarity readily gives moles. The same relationship also allows us to find the number of moles of solute present if we know both the volume and molarity of any solution. If we use n to represent the number of moles, V for volume, and M for molarity, we can simply rearrange the definition of molarity: n=M×V


Molarity is another ratio between otherwise unrelated variables—similar to mass density. It can be used in calculations in the same way as density.

(3.2)

3 .10

Aqueous solutions of sodium hypochlorite (NaClO) can be used in the synthesis of hydrazine (N2H4). Hydrazine has often been used as a rocket fuel, and a derivative of hydrazine is used as a fuel for engines in the orbital maneuvering system of the space shuttle. A solution is prepared by dissolving 45.0 g of NaClO in enough water to produce exactly 750 mL of solution. What is the molarity of the solution?

Strategy To obtain molarity we need two quantities—moles of solute and liters of solution. Neither of these quantities is given directly in the question, but both can be obtained readily from information given. We must use the molar mass of NaClO to convert from mass to moles and convert the volume from milliliters to liters. Then we can use the definition of molarity to obtain the molarity of the solution.

Hydrazine, N2H4

Solution First we compute the moles of solute. 1.00 mol NaClO 45.0 g NaClO × ———————— = 0.604 mol NaClO 74.442 g NaClO Then convert the solution volume from mL to L: 1.00 L 750 mL × ————— = 0.750 L 1000 mL Finally, calculate the molarity: 0.604 mol moles of solute Molarity = ———————— = ————— = 0.806 M NaClO 0.750 L liters of solution

Analyze Your Answer Our answer is somewhat smaller than, but fairly close to, 1 M. Does this make sense? The amount of NaClO was about 2/3 of a mole and the volume of the solution was 3/4 of a liter. The ratio of 2/3 to 3/4 is 8/9, close to but a little bit smaller than 1, so our result seems reasonable.

Discussion As usual, this problem could also be set up as a single calculation including the same conversions. Whichever approach you choose, you should always check to be sure that your units work out correctly. Here we have units of mol/L for our answer, which is appropriate for molarity. Confirming this helps avoid careless errors.

Check Your Understanding To produce hydrazine, sodium hypochlorite must react with ammonia. How many moles of NH3 would be in a 4.0-L bottle of 14.8 M ammonia solution?


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Dilution One of the most common procedures encountered in any science laboratory is dilution, the process in which solvent is added to a solution to decrease the concentration of the solute. In dilution, the amount of solute does not change. The number of moles of solute is the same before and after dilution. Because we already know that the number of moles of solute equals the product of molarity and volume, we can write the following equation, where the subscripts denote the initial (i) and final (f) values of the quantities involved: This dilution formula can be used when there is only one solute present. If a reaction is taking place, the problem must be approached as a stoichiometry problem, which we will take up in Chapter 4.

Mi × Vi = Mf × Vf

(3.3)

The simplicity of equations such as this is alluring. One of the most common errors made by students in chemistry is to use this relationship in situations where it is not valid. The only place to use this relationship is in the process of dilution or concentration of a solution. Example Problem 3.11 illustrates a common laboratory situation where this equation is valid and useful.

E XAM PL E PRO BLEM

3 .11

A chemist requires 1.5 M hydrochloric acid, HCl, for a series of reactions. The only solution available is 6.0 M HCl. What volume of 6.0 M HCl must be diluted to obtain 5.0 L of 1.5 M HCl?

Strategy First we need to recognize that this is the dilution of the concentrated solution to prepare the desired one. The underlying concept we must use is that the number of moles of HCl will be the same before and after dilution, and this means we can use Equation 3.3. We know the desired final molarity and final volume, as well as the initial molarity. So we can solve for the needed initial volume.

Solution Initial concentration of HCl: Mi = 6.0 M Final concentration of HCl: Mf = 1.5 M Final volume of solution: Vf = 5.0 L The unknown is the initial volume, Vi. Rearranging Equation 3.3, Mf × Vf Vi = ———— Mi Inserting the known quantities on the right-hand side, 1.5 M × 5.0 L Vi = ——————— = 1.3 L 6.0 M To obtain the desired quantity of diluted HCl, the chemist should begin with 1.3 L of the concentrated solution and add enough water to bring the volume up to 5.0 L.

Analyze Your Answer Once we are certain that we are dealing with a dilution problem, one way to make sure the answer makes sense is to think about which solution is concentrated and which is dilute. You will always need smaller volumes of the concentrated solution. Did we get it right this time? It seems likely that we did: we were asked to calculate how much of the concentrated solution we would need, so the volume we calculate should be the smaller of the two volumes involved in the problem.

Check Your Understanding If 2.70 mL of 12.0 M NaOH is diluted to a volume of 150.0 mL, what is the final concentration?

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INSIGHT IN TO

3.6

Explosives and Green Chemistry

Concern for the environment and the impact of human activities on the environment has grown dramatically in recent years. This concern led to increased scrutiny of the chemical industry, which is often viewed as a significant source of pollution. The chemical industry has responded by developing the concept of green chemistry, which refers to the philosophy that chemical processes and products should be designed with the goal of reducing their environmental impacts. The U.S. Environmental Protection Agency (EPA) has endorsed a set of twelve principles to guide the implementation of green chemistry practices. These principles encompass such broad concepts as maximizing the extent to which the raw materials end up in the final product, eliminating or minimizing the use of toxic solvents and reactants, improving energy efficiency, and minimizing the production of any waste materials as by-products. Various programs, such as the Presidential Green Chemistry Challenge Awards, have been created to encourage both industrial and academic chemists to pursue these principles of green chemistry. The idea of a “green explosive” may seem rather unlikely at first. But, in actuality, the explosives industry has begun to examine the environmental impact of many of its practices and is now seeking to develop greener alternatives. One of the most common uses of explosives is in firearms. The firing of a gun is usually a two-step sequence of small explosions. First, either the hammer of the gun or an electric current produced by the firing mechanism detonates a “primer” explosive. This, in turn, induces a larger explosion that actually propels the bullet out of the barrel of the gun. For many years the most common primers have been lead-based substances such as lead azide (Pb(N3)2). But the toxic lead is left behind after the explosion of the primer. Although the amount of lead involved in a single firing is quite small, the cumulative effects are significant: heavily used firing ranges, such as those used by the military or police, frequently show levels of lead that are considered dangerously high. Removing the lead from such sites poses serious challenges. (Figure 3.10 shows the type of equipment often used for this type of environmental remediation.) Research is underway to find replacement primers that would lessen this environmental impact. One leading candidate uses nanoparticles of aluminum, which react violently with oxygen. Larger scale explosives are also coming under scrutiny. Many military explosives are stored for long periods of time and eventually destroyed rather than detonated. Such disposal usually involves a controlled burn, which often generates atmospheric

Environmental chemistry is the study of chemical phenomena in the environment. Green chemistry is an attempt to reduce or eliminate the environmental impact of chemical processes.

Figure 3.10 ❚ Gravity separators

Image courtesy of www.extrac-tec.com

like this one from Extrac-tec are used in remediation of lead from firing ranges. Differences in density and solubility allow the equipment to extract lead or other heavy metals from soil.

3.6 Explosives and Green Chemistry

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pollutants such as carbon monoxide. This has led to a push to develop explosives in which the carbon content is minimized. (Needless to say, such research can be extremely dangerous and requires stringent safety measures.) As you read through this chapter, you most likely have noticed how many of the explosives we have mentioned contain nitrogen atoms, which then form N2 as a product of the explosion. This is because the formation of the nitrogen-nitrogen bond in N2 releases a great deal of energy. So the search for explosives that contain less carbon has focused largely on nitrogen-rich compounds. Early in 2009, a group of German chemists reported the synthesis of a number of compounds containing the azidotetrazolate ion (CN72). One example is hydrazinium azidotetrazolate, with the formula N2H5CN7. The low carbon content of such a molecule means that even when exploded, this material would generate very little CO2 or CO, thus reducing the environmental impact. It also means that the accumulation of carbon-based by-products (soot, etc.) in gun or artillery barrels would be significantly reduced. As with many applications of green chemistry, conversion to these new explosives is likely to increase some costs in the short term. But the need to include longer-term environmental impacts as a serious consideration is increasingly recognized. There is very little doubt that the principles of green chemistry will continue to grow in importance for the foreseeable future. In addition to minimizing future environmental impacts, there is also a compelling need to develop ways to remediate existing environmental pollution, such as the accumulated lead from firing ranges. Keeping that lead from migrating into the water system is of particular importance. One possible strategy involves the limited solubility of many lead compounds. Lead phosphate, for example, is not soluble in water. So adding phosphates to contaminated soil in firing ranges can induce precipitation reactions, and the resulting lead phosphate product can be trapped and removed. You may be aware that phosphates themselves are often viewed as pollutants, so this example illustrates the complexity of the chemistry involved in many environmental issues.

FOCU S O N

PRO BLEM S O LVI NG

Question When heated in air, some element, X, reacts with oxygen to form an oxide with the chemical formula X2O3. If a 0.5386-g sample of this unknown element yields an oxide with a mass of 0.7111 g, describe how you would determine the atomic mass of the element.

Strategy Chemical formulas are expressed in terms of numbers of particles. In this problem, we have to start with masses and somehow infer information about the numbers of particles. We did a similar problem in this chapter: the determination of an empirical formula from percentage by mass data. In this case, from the experiment described, we can determine the mass of each element in the unknown compound— providing the same type of data we had from mass percentage calculations. Then, because we know the empirical formula and we know the molar mass of oxygen, we can determine the molar mass of the unknown element.

Solution First we would determine the mass composition. The mass of element X in the oxide sample must be assumed to be the 0.5386 g that was used to generate the oxide. Subtracting this value from the 0.7111 g of the oxide sample then provides the mass of the oxygen. The mass of oxygen can be converted to moles of oxygen using the molar mass (which could be looked up if we don’t know it). Then, we could use the mole ratio from the empirical formula given in the problem to assert that whatever value we obtain, the number of moles of X would be two-thirds of this result. (This follows from the fact that there are two X atoms for every three oxygen atoms.) Then, we would have a value for the number of moles of X and the mass of X in the sample—taking the ratio provides the molar mass, from which we could use a periodic table to look up the identity of element X.

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S U M M ARY Chemical reactions are the most important events in chemical systems. So the ability to describe reactions succinctly is vital to our study of chemistry. Such descriptions rely on the balanced chemical equation, a written representation of the reaction accurately reflecting the fact that matter is neither created nor destroyed. One requirement for a reaction to occur is that the chemicals involved must be able to mix and interact with one another. Thus the physical state of the reactants can be important. In particular, an aqueous solution is often used as the medium for a reaction. To describe reactions in the aqueous phase effectively, we often need to specify the molar concentration of the solutions used. We also have options about how to write the chemical reaction itself—as a molecular equation, for example, or a net ionic equation. Several classes of chemical reactions, such as precipitation reactions or acid–base neutralizations, are sufficiently common that additional definitions have been established to improve our ability to communicate about them. Ultimately, numerical information is often desirable, so chemists have devised a way to count the enormous numbers of

particles involved in reactions on laboratory or industrial scales. The unit chemists use to count atoms or molecules is called the mole, and it is specified by Avogadro’s number, 6.022 × 1023. The molar mass of an element can be obtained readily from the periodic table, and that of a compound can be found by adding the molar masses of its constituent atoms. Because we can determine the molar mass of virtually any chemical substance, we can convert between a mass (that can be readily measured) and the number of particles in moles (that can be used in chemical calculations.) Molar concentrations provide a similar ability to convert between a measurable quantity (in this case volume of solution) and moles. There are several types of calculations that can be undertaken once we recognize the concept that a mole is simply a way to count particles. We can determine the empirical formula of a compound based on the mass percentage analysis of its elements. We can also determine the molar concentration of solutions prepared by dissolving a known mass of the solute or by diluting a concentrated solution.

K EY TE RM S acid (3.3)

law of conservation of matter (3.2)

salt (3.3)

aqueous solution (3.3)

molar concentration (3.5)

solubility (3.3)

Avogadro’s number (3.4)

molar mass (3.4)

soluble (3.3)

base (3.3)

molarity (3.5)

solute (3.3)

chemical equation (3.2)

mole (3.4)

solution (3.3)

concentration (3.3)

mole ratio (3.4)

solvent (3.3)

dilution (3.5)

molecular equation (3.3)

spectator ions (3.3)

dissociation reaction (3.3)

net ionic equation (3.3)

stoichiometric coefficients (3.2)

electrolyte (3.3)

neutralization (3.3)

stoichiometry (3.2)

elemental analysis (3.5)

nonelectrolyte (3.3)

strong electrolyte (3.3)

explosives (3.1)

photochemical reaction (3.2)

total ionic equation (3.3)

fuel cell (3.6)

precipitation reaction (3.3)

weak electrolyte (3.3)

green chemistry (3.6)

products (3.2)

insoluble (3.3)

reactants (3.2)



INSIGHT INTO Explosives 3.1 List at least three traits shared by most explosives. 3.2 How does nitroglycerin differ from most modern explosives? Use the web to determine if any explosives based on nitroglycerin are used today. 3.3 In the chapter, it was noted that a blast front forms when an explosive is used. Use the web to find more information about how a blast front propagates and draw a diagram that depicts this information.

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3.4 How is an explosive similar to a fuel? How do explosives and fuels differ? 3.5 Many explosives contain nitrogen atoms. Why do you think this is so? Chemical Formulas and Equations 3.6 Define the terms reactants and products as they refer to chemical reactions. 3.7 Which symbols are used to indicate solids, liquids, gases, and aqueous solutions in chemical equations?


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3.8 How is the addition of heat symbolized in a chemical equation? The addition of light energy?

(b) production of methanol, CH3OH(ℓ), by combining H2(g) and CO(g)

3.9 What law of nature underpins the concept of a balanced chemical equation?

(c) production of sulfuric acid by combining sulfur, oxygen, and water

3.10 Define the term stoichiometric coefficient. 3.11

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Balance these equations.

(a) Al(s) + O2(g) : Al2O3(s) (b) N2(g) + H2(g) : NH3(g) (c) C6H6(ℓ) + O2(g) : H2O(ℓ) + CO2(g) 3.12

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Balance the following equations.

(a) CaC2(s) + H2O(ℓ) : Ca(OH)2(s) + C2H2(g) (b) (NH4)2Cr2O7(s) : Cr2O3(s) + N2(g) + H2O(g) (c) CH3NH2(g) + O2(g) : CO2(g) + N2(g) +H2O(g) 3.13 An explosive whose chemical formula is C3H6N6O6 produces water, carbon dioxide, and nitrogen gas when detonated in oxygen. Write the chemical equation for the detonation reaction of this explosive. 3.14 A number of compounds are used in cement and reactions among them occur when water is added. In one, CaO reacts with Al2O3 and water to form Ca3Al2(OH)12. Write a balanced chemical equation for this process.

3.18 Diborane and related compounds were proposed as rocket fuels in the 1950s. A representative reaction for this class of molecules is that of B2H6 and O2 to form B2O3 and H2O. Write the balanced chemical equation for this process. 3.19 Silicon nitride, Si3N4, is used as a reinforcing fiber in construction materials. It can be synthesized from silicon tetrachloride and ammonia. The other product is ammonium chloride. Write the balanced chemical reaction for this process. 3.20 The pictures below show a molecular scale view of a chemical reaction between the compounds AB2 and B2. (A atoms are shown in blue and B atoms in white). The box on the left represents the reactants at the instant of mixing, and the box on the right shows what is left once the reaction has gone to completion. Write a balanced chemical equation for this reaction. As usual, your equation should use the smallest possible whole number coefficients for all substances.

3.15 Ethanol, C 2 H 5 OH, is found in gasoline blends used in many parts of North America. Write a balanced chemical equation for the combustion of C2H5OH to form CO2 and H2O.

Aqueous Solutions and Net Ionic Equations 3.21 Define the terms solution, solute, and solvent. Ethanol, C2H5OH

3.16

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Balance the following equations.

(a) reaction to produce “superphosphate” fertilizer:

3.22 When a substance dissolves, does a chemical reaction always occur? What happens if it is not a reaction? 3.23 What is a concentrated solution? A dilute solution? 3.24 Define the term saturated solution. 3.25

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3.26

NaBH4(s) + H2SO4(aq) : B2H6(g) + H2(g) + Na2SO4(aq) (c) reaction to produce tungsten metal from tungsten(VI) oxide:

■ Classify the following compounds as electrolytes or non-electrolytes: (a) potassium chloride, KCl, (b) hydrogen peroxide, H2O2, (c) methane, CH4, (d) barium nitrate, Ba(NO3)2.

3.27

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3.28

■ Decide whether each of the following is water-soluble. If soluble, tell what ions are produced. (a) Na 2 CO 3 , (b) CuSO4, (c) NiS, (d) BaBr2

WO3(s) + H2(g) : W(s) + H2O(ℓ) (d) decomposition of ammonium dichromate: (NH4)2Cr2O7(s) : N2(g) + H2O(ℓ) + Cr2O3(s) 3.17

Write balanced chemical equations for the following reactions. ■

(a) production of ammonia, NH3(g), by combining N2(g) and H2(g)

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What is an electrolyte? How can you differentiate experimentally between a weak electrolyte and a strong electrolyte? Give an example of each.

Ca3(PO4)2(s) + H2SO4(aq) : Ca(H2PO4)2(aq) + CaSO4(s) (b) reaction to produce diborane, B2H6:


The following compounds are water-soluble. What ions are produced by each compound in aqueous solution? (a) KOH, (b) K2SO4, (c) LiNO3, (d) (NH4)2SO4

3.29 The solubility of NaCl in water is 35.7 g NaCl/100 g H2O. Suppose that you have 500.0 g of NaCl. What is the minimum volume of water you would need to dissolve it all? (Assume that the density of water is 1.0 g/mL.) 3.30

A packaging engineer is working on a new design for cold packs. A burst chamber in the pack will contain

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14.0 mL of water, which will be released and used to dissolve ammonium nitrate. (a) If the solubility of NH4NO3 is 190 g/100 g H2O, what mass of this compound should be used with this amount of water? (b) The design specifications indicate that these chemicals can comprise no more than 35% of the mass of the product. What would the product weigh if it just meets these design specs? 3.31

Classify each of these as an acid or a base. Which are strong and which are weak? What ions are produced when each is dissolved in water? (a) KOH, (b) Mg(OH)2, (c) HClO, (d) HBr, (e) LiOH, (f ) H2SO3 ■

sulfide, galena, (b) C 2H 6, ethane, a hydrocarbon fuel, (c) CH3COOH, acetic acid, an important ingredient in vinegar, (d) NH4NO3, ammonium nitrate, a fertilizer 3.44 When an automobile air bag deploys, solid sodium azide decomposes rapidly to produce nitrogen gas, and this gas inflates the bag. What is the molar mass of sodium azide (NaN3)? 3.45 Calculate the molar mass of the following compounds: (a) magnesium phosphate, (b) sodium sulfide, (c) dinitrogen tetroxide

3.32 Define the term spectator ion.

3.46 Determine the molar mass of these ceramic materials. (a) HfN, (b) ThO2, (c) BaTiO3

3.33 What is the difference between a total ionic equation and a net ionic equation?


3.34

Balance the following equations and then write the net ionic equation.

3.47 A chemist needs exactly 2 moles of NaOH to make a solution. What mass of NaOH must be used?

(a) (NH4)2CO3(aq) + Cu(NO3)2(aq) : CuCO3(s) + NH4NO3(aq)

3.48 What mass of ozone (O 3 ) gives 4.5 moles of the substance?

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(b) Pb(OH)2(s) + HCl(aq) : PbCl2(s) +H2O(,) (c) BaCO3(s) + HCl(aq) : BaCl2(aq) + H2O(,) + CO2(g) (d) CH3CO2H(aq) + Ni(OH)2(s) : Ni(CH3CO2)2(aq) + H2O(,) 3.35

Balance the following equations and then write the net ionic equation.

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Ozone, O3

(a) Zn(s) + HCl(aq) : H2(g) + ZnCl2(aq) (b) Mg(OH)2(s) + HCl(aq) : MgCl2(aq) + H2O(,) (c) HNO3(aq) + CaCO3(s) : Ca(NO3)2(aq) + H2O(,) + CO2(g) (d) (NH4)2S(aq) + FeCl3(aq) : NH4Cl(aq) + Fe2S3(s) 3.36 In principle, it may be possible to engineer the trapping of carbon dioxide by a precipitation reaction with Ca(OH)2 to form calcium carbonate and water. Write the molecular, total ionic, and net ionic equations for this reaction.

3.49

■ Calculate the mass in grams of each the following: (a) 2.5 mol of aluminum, (b) 1.25 × 10−3 mol of iron, (c) 0.015 mol of calcium, (d) 653 mol of neon

3.50

Calculate the mass in grams of 13.5 mol of (a) vinyl chloride, C 2 H 3 Cl, the starting material for a plastic, (b) capsaicin, C18H27NO3, the substance that makes red chili peppers “hot,” and (c) stearic acid, C18H36O2, used in soaps.

3.51

How many moles are present in the given quantities of explosives? (a) 358.1g trinitrotoluene (TNT), C7H5N3O6, (b) 82.6 g nitromethane, CH 3NO 2, (c) 1.68 kg RDX, C3H6N6O6

3.52

A test of an automobile engine’s exhaust revealed that 3.7 g of NO2 was emitted in 10 minutes of operation. How many moles of NO2 would this engine release if it were used for a 45-minute commute, assuming that this measured number is representative of the emission under all circumstances?

Interpreting Equations and the Mole 3.37 Explain the concept of the mole in your own words. 3.38 How many entities are present in (a) 3.21 mol argon atoms (b) 1.3 × 10−12 mol marbles (c) 3 × 106 mol water molecules (d) 7.63 ×10−18 mol basketballs 3.39 If a typical grain of sand occupies a volume of 1.3 × 10−4 cm3, what is the volume (in cm3) of one mole of sand (ignoring the space between grains)? What is the volume in liters? 3.40 Estimate the size of a particle one mole of which would fill your bedroom or dorm room. Calculate the molar mass of each of the following compounds. (a) Fe2O3, iron(III) oxide, (b) BCl3, boron trichloride, (c) C6H8O6, ascorbic acid (vitamin C).

3.41

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3.42

Calculate the molar masses (in grams per mole) of (a) cane sugar, C 12H 22O 11, (b) laughing gas, N 2O, and (c) vitamin A, C20H30O.

3.43

■ Calculate the molar mass of each of these compounds and the mass percentage of each element. (a) PbS, lead(II)

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3.53 Modern instruments can measure a mass as small as 5 nanograms. If one observed 5.0 ng of CO2, how many molecules were measured? 3.54 How many H atoms are present in 7.52 g of propane, C3H8?

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Propane, C3H8


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3.55 How many N atoms are present in 214 g of TNT, C7H5N3O6?

3.67 Determine the final molarity for the following dilutions: (a) 24.5 mL of 3.0 M solution diluted to 100.0 mL

3.56 A sample of H2C2O4 ? 2H2O of mass 3.35 g is heated to drive off the waters of hydration (designated separately in the chemical formula). What mass of H2C2O4 remains once the water has been removed by heating? 3.57 An average person inhales roughly 2.5 g of O2 in a minute. How many molecules of oxygen are inhaled in (a) one minute, (b) one hour, (c) one day by an average person? 3.58

3.60

■ Determine the simplest formulas of the following compounds: (a) the food enhancer monosodium glutamate (MSG), which has the composition 35.51% C, 4.77% H, 37.85% O, 8.29% N, and 13.60% Na, (b) zircon, a diamond-like mineral, which has the composition 34.91% O, 15.32% Si, and 49.76% Zr, (c) nicotine, which has the composition 74.0% C, 8.65% H, and 17.4% N

3.61

■ The composition of materials such as alloys can be described in terms of mole percentage (mol %), atom percentage (at %), or weight percentage (wt %). Carry out the following conversions among these units. (a) 60 wt % Cu and 40 wt % Al to at %, (b) 25 mol % NiO and 75 mol % MgO to wt %, (c) 40 wt % MgO and 60 wt % FeO to mol %

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Copper can have improved wear resistance if alloyed with ceramic alumina, Al2O3. If a copper alloy has 8.5 wt % Al2O3, what is its composition in mol %?

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3.63 Calculate the molarity of each of the following solutions: (a) 1.45 mol HCl in 250. mL of solution (b) 14.3 mol NaOH in 3.4 L of solution (c) 0.341 mol KCl in 100.0 mL of solution (d) 250 mol NaNO3 in 350 L of solution ■ What is the molarity of each ion present in aqueous solutions prepared by dissolving 20.00g of the following compounds in water to make 4.50 L of solution? (a) cobalt(III) chloride, (b) nickel(III) sulfate, (c) sodium permanganate, (d) iron(II) bromide

3.65 How many moles of solute are present in each of these solutions? (a) 48.0 mL of 3.4 M H2SO4 (b) 1.43 mL of 5.8 M KNO3 (c) 3.21 L of 0.034 M NH3 (d) 1.9 × 10 L of 1.4 × 10 3

−5

M NaF

3.66 How many grams of solute are present in each of these solutions? (a) 37.2 mL of 0.471 M HBr (b) 113.0 L of 1.43 M Na2CO3 (c) 21.2 mL of 6.8 M CH3COOH (d) 1.3 × 10−4 L of 1.03 M H2SO3

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(d) 2.35 L of 12.5 M solution diluted to 100.0 L 3.68 Determine the initial volume (in mL) needed to generate the desired solutions by dilution: (a) 10.0 L of 0.45 M solution from 3.0 M solution

A large family of boron-hydrogen compounds has the general formula BxHy. One member of this family contains 88.5% B; the remainder is hydrogen. What is its empirical formula? Mandelic acid is an organic acid composed of carbon (63.15%), hydrogen (5.30%), and oxygen (31.55%). Its molar mass is 152.14 g/mol. Determine the empirical and molecular formulas of the acid.

3.64

(c) 1.45 mL of 0.034 M solution diluted to 10.0 mL

(b) 1.50 L of 1.00 M solution from 11.7 M solution

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3.59

3.62

(b) 15.3 mL of 4.22 M solution diluted to 1.00 L


(c) 100.0 mL of 0.025 M solution from 1.15 M solution (d) 50.0 mL of 3.3 × 10−4 M solution from 0.25 M solution 3.69

Commercially available concentrated sulfuric acid is 18.0 M H 2SO 4. Calculate the volume of concentrated sulfuric acid required to prepare 2.50 L of 1.50 M H2SO4 solution.

3.70

■ Magnesium is lighter than other structural metals, so it is increasingly important in the design of more efficient vehicles. Mg2+ ions are present in seawater, and the metal is often prepared by “harvesting” these ions and converting them to Mg metal. The average magnesium content of the oceans is about 1270 g Mg2+ per ton of seawater, and the density of seawater is about 1.03 g/mL. What is the molarity of Mg2+ ions in seawater? The design for a concept car calls for 103 kilograms of magnesium per vehicle. How many gallons of seawater would be required to supply enough magnesium to build one of these cars?

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INSIGHT INTO Explosives and Green Chemistry 3.71 Lead based primers were introduced years ago to replace toxic mercury(II)fulminate. If a percussion cap contains 0.550 g of Hg(CNO)2, how many atoms of mercury are present? 3.72

■ One lead-based primer is lead styphnate, C6HN3O8Pb. (a) What is the molar mass of lead styphnate? (b) What is the percentage by mass of lead in lead styphnate?

3.73 Diazodinitrophenol, C6H2N4O5, is a component of some recently patented “green” primers. Assume the combustion products of the reaction of this compound with oxygen are CO2, H2O, and N2 and write the balanced chemical equation for the reaction. 3.74 Two new candidates for green primers are called HBT (C 2H 4N 10) and G2Z (C6H 14N 22). (a) Which compound has a higher percentage of nitrogen by mass? (b) Calculate the molar mass of each compound. (c) If the detonation of both compounds in oxygen results in the production of CO2, H2O, and N2, write balanced chemical equations for these reactions. 3.75 Use the web to research environmental remediation projects for firing ranges. What is the most important source of pollutants in the site you choose? Is the introduction of green explosive primers a solution to that specific problem? 3.76 Use the web to identify characteristics required for a material to be useful as a primer in bullet manufacturing. Which of these characteristics are the most challenging for green explosives to achieve?

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unbalanced net ionic equation for this reaction is given below. Balance it.

Additional Problems 3.77 Nitric acid (HNO3) can be produced by the reaction of nitrogen dioxide (NO2) and water. Nitric oxide (NO) is also formed as a product. Write a balanced chemical equation for this reaction.

CaCO3(s) + H3O+(aq) : H2O(ℓ) + CO2(g) + Ca2+(aq) 3.85

Answer each of the following questions. Note that none of them should require difficult calculations.

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(a) How many molecules are present in 1 mole of octane (C8H18)? (b) How many moles of fluorine atoms are present in 4 moles of C2F6? (c) What is the approximate mass (to the nearest gram) of 3 × 1023 carbon atoms? 3.86

Nitric acid, HNO3

3.78 One step in the enrichment of uranium for use in nuclear power plants involves the reaction of UO2 with hydrofluoric acid (HF) solution. The products are solid UF4 and water. Write a balanced chemical equation for this reaction. 3.79 Pyridine has the molecular formula C5H5N. When pyridine reacts with O2, the products are CO2, H2O, and N2. Write a balanced equation for this reaction. 3.80 Pyrrole has the molecular formula C4H5N. When pyrrole reacts with O2, the products are CO2, H2O, and N2. Write a balanced equation for this reaction. 3.81 Hydrogen cyanide (HCN) is extremely toxic, but it is used in the production of several important plastics. In the most common method for producing HCN, ammonia (NH3) and methane (CH4) react with oxygen (O2) to give HCN and water. Write a balanced chemical equation for this reaction.

Hydrogen cyanide, HCN

3.82 Many chemical reactions take place in the catalytic converter of a car. In one of these reactions, nitric oxide (NO) reacts with ammonia (NH3) to give nitrogen (N2) and water. Write a balanced equation for this reaction. 3.83 Adipic acid is used in the production of nylon, so it is manufactured in large quantities. The most common method for the preparation of adipic acid is the reaction of cyclohexane with oxygen. Balance the skeleton equation shown below.

(a) Which sample contains more atoms? (b) Which sample contains more molecules? (c) Which sample has a larger mass? 3.87 Cumene is a hydrocarbon, meaning that it contains only carbon and hydrogen. If this compound is 89.94% C by mass and its molar mass is 120.2 g/mol, what is its molecular formula? 3.88 Methyl cyanoacrylate is the chemical name for the substance sold as Super Glue, and it has the chemical formula C5H5NO2. Calculate the number of molecules of this substance in a 1.0-ounce tube of Super Glue, assuming that the glue is 80% methyl cyanoacrylate by mass. 3.89 A low-grade form of iron ore is called taconite and the iron in the ore is in the form Fe3O4. If a 2.0-ton sample of taconite pellets yields 1075 pounds of iron when it is refined, what is the mass percentage of Fe3O4 in taconite? 3.90 The characteristic odor of decaying flesh is due to the presence of various nitrogen-containing compounds. One such compound, called putrescine, was analyzed and found to contain 54.49% carbon, 13.72% hydrogen, and 31.78% nitrogen by mass. If the molar mass of putrescine is known to be between 85 and 105, what is its molecular formula? 3.91 Iron-platinum alloys may be useful as high-density recording materials because of their magnetic properties. These alloys have been made with a wide range of composition, from 34.0 at % Pt to 81.8 at % Pt. Express this range in mol %. 3.92

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3.93

■ Which (if any) of the following compounds are electrolytes? (a) glucose, C 6H 12O 6, (b) ethanol, C 2H 5OH, (c) magnesium sulfide, MgS, (d) sulfur hexafluoride, SF6

3.94

■ Classify the following compounds as acids or bases, weak or strong. (a) perchloric acid, (b) cesium hydroxide, (c) carbonic acid, H2CO3, (d) ethylamine, C2H5NH2

O ⫹ O2

OH

HO

⫹ H2O

O cyclohexane

adipic acid

3.84 Calcium carbonate (limestone, CaCO3) dissolves in hydrochloric acid, producing water and carbon dioxide. An

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■ Consider two samples of liquid: one mole of water (H2O) and one mole of ethanol (C2H5OH). Answer each of the following questions. Note that none of these should require significant calculations.

Some aluminum-lithium alloys display the property of superplasticity, meaning they can undergo tensile deformation by large amounts (1000 times or more) without breaking. If such an alloy has 4 wt % Li, what is its composition in mol %? Explain the relative magnitudes of the mole percentage and weight percentage based on the molar masses of aluminum and lithium.


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3.95

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3.96

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3.97

■ Nitric acid is often sold and transported as a concentrated 16 M aqueous solution. How many gallons of such a solution would be needed to contain the entire 1.65 × 107 pounds of HNO3 produced in the United States in 2008?

3.98

3.99

3.100

3.101

What is the mass in grams of solute in 250. mL of a 0.0125 M solution of KMnO4?

What volume of 0.123 M NaOH in milliliters contains 25.0 g NaOH?

Twenty-five mL of a 0.388 M solution of Na2SO4 is mixed with 35.3 mL of 0.229 M Na2SO4. What is the molarity of the resulting solution? Assume that the volumes are additive. ■

■ As computer processor speeds increase, it is necessary for engineers to increase the number of circuit elements packed into a given area. Individual circuit elements are often connected using very small copper “wires” deposited directly onto the surface of the chip. In current generation processors, these copper interconnects are about 65 nm wide. How many copper atoms would be in a 1-mm length of such an interconnect, assuming a square cross section. (The density of copper is 8.96 g/cm3.) ■ As chip speeds increase, the width of the interconnects described in Problem 3.99 must be reduced. A hypothetical limit to this process would be reached if the interconnect was just one copper atom wide. Use the density of copper (8.96 g/cm3) to estimate the diameter of a copper atom. (Optical interconnects are being developed, and are likely to replace copper in this application within a few years.)

Materials engineers often create new alloys in an effort to improve the properties of an existing material. ZnO based semiconductors show promise in applications like light-emitting diodes, but their performance can be enhanced by the addition of small amounts of cadmium. One material that has been studied can be represented by the formula Zn0.843Cd0.157O. (These materials are solid solutions, and so they can have variable compositions. The non-integer coefficients do not imply fractional atoms.) Express the composition of this alloy in terms of (a) at %, (b) mol %, and (c) wt %. ■

FOCUS ON PROBLEM SOLVING EXERCISES 3.102 The protein that carries oxygen in the blood is called hemoglobin. It is 0.335% Fe by mass. Given that a molecule of hemoglobin contains four iron atoms, describe how you can calculate the molar mass of hemoglobin. Do you need to look up anything to do it? 3.103 The chlorophyll molecule responsible for photosynthesis in plants has 2.72% Mg by mass. There is only one Mg atom per chlorophyll molecule. How can you determine the molar mass of chlorophyll based on this information? 3.104 In one experiment, the burning of 0.614 g of sulfur produced 1.246 g of sulfur dioxide as its only product. In a second experiment run to verify the outcome of the first,

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a sample was burned and 0.813 g of sulfur dioxide was obtained, but the student running the experiment forgot to determine the initial mass of sulfur burned. Describe how the data already obtained could be used to calculate the initial mass of the sulfur burned in the second experiment. 3.105 MgCl2 is often found as an impurity in table salt (NaCl). If a 0.05200-g sample of table salt is found to contain 61.10% Cl by mass, describe how you could determine the percentage of MgCl2 in the sample. 3.106 Under the Clean Air Act, companies can purchase the right to emit certain pollutants from other companies that emit less than they are allowed by law. Suppose that the cost to purchase a pollution credit for 1 ton of sulfur dioxide is $300. Describe how you would determine how many SO2 molecules could be emitted per dollar at this price. 3.107 The average person exhales 1.0 kg of carbon dioxide in a day. Describe how you would estimate the number of CO2 molecules exhaled per breath for this average person. 3.108 The simplest approximate chemical formula for the human body could be written as C728H4850O1970N104Ca24P16 K4S4Na3Cl2Mg. Based on this formula, describe how you would rank by mass the 10 most abundant elements in the human body. 3.109 For the oxides of iron, FeO, Fe2O3, and Fe3O4, describe how you would determine which has the greatest percentage by mass of oxygen. Would you need to look up any information to solve this problem? Cumulative Problems 3.110 Consider common explosives such as TNT (C7H5N3O6) and halite (C2H6N4O4). What type of chemical bonding would you expect to find in these chemicals? 3.111 For the compounds that commonly undergo precipitation reactions, what type of chemical bonding is expected? 3.112 If you have 32.6 g of sodium carbonate that is dissolved to give 2.10 L of solution, what is the molarity of the solution? What is the molarity of the sodium ions? 3.113 If you have 21.1 g of iron(II) nitrate that is dissolved to give 1.54 L of solution, what is the molarity of the solution? What is the molarity of the nitrate ions? 3.114 What type of reasoning were we using when we developed the equation for dilution, Mi × Vi = Mf × Vf? 3.115 Most periodic tables provide molar masses with four or five significant figures for the elements. How accurately would you have to measure the mass of a sample of roughly 100 g to make a calculation of the number of moles of the chemical to have its significant figures limited by the molar mass calculation rather than the mass measurement?

■


4

Stoichiometry

OUTLINE 4.1

INSIGHT INTO Gasoline and

Other Fuels 4.2

Fundamentals of Stoichiometry

4.3

Limiting Reactants

4.4

Theoretical and Percentage Yields

4.5

Solution Stoichiometry

4.6

INSIGHT INTO Alternative Fuels and Fuel Additives

Chemical reactions initiated in the Pegasus rocket accelerate NASA’s X-43A aircraft to hypersonic speeds. In a test flight, the X-43A approached Mach 10, ten times the speed of sound. NASA Dryden Flight Research Center

A

tremendous number of chemical compounds exist in nature and undergo myriad reactions. By building and exploiting a systematic understanding of reactivity, chemists also have produced an impressive array of man-made compounds. Many of the explosives and polymers that we’ve encountered in the last two chapters are handy examples, as are most modern pharmaceuticals. Other practical applications of chemistry involve the development of new routes to synthesize existing natural compounds. As you might imagine, making any of these syntheses commercially feasible requires detailed and quantitative understanding of the reactions involved. The economics of any chemical process obviously depend on the amounts of each reactant needed to produce a given amount of product. For processes carried out on an industrial scale, even very small changes in efficiency can have enormous impact on profitability. The quantitative relationships between the amounts of reactants and products in a chemical reaction are referred to as stoichiometry. In this chapter, we’ll use a variety of chemical reactions associated with various fuels to illustrate important stoichiometric concepts.

The word stoichiometry derives from two Greek words: stoicheion (meaning “element”) and metron (meaning “measure”).

Chapter Objectives After mastering this chapter, you should be able to ❚ describe the chemical composition of gasoline. ❚ write balanced chemical equations for the combustion of fuels.

Online homework for this chapter may be assigned in OWL.

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❚ calculate the amount of product expected from a chemical reaction, given the amounts of reactants used. ❚ calculate the amounts of reactants needed in a chemical reaction to produce a specified amount of product. ❚ identify a limiting reactant and calculate the amount of product formed from a nonstoichiometric mixture of reactants. ❚ calculate the percentage yield of a chemical reaction. ❚ identify at least two common additives in gasoline and explain why they are used.

INSIGHT INTO

4.1

Gasoline and Other Fuels

Gasoline is the most widely used fuel in our society: daily consumption in the United States averages more than 350 million gallons. To explore the chemistry of gasoline, we must begin by examining its composition. The fuel we know as gasoline is actually a rather complex mixture, typically containing more than 100 different chemical compounds. The exact composition of gasoline varies somewhat depending on factors including the grade of gas, geographic location, and time of year. But the predominant compounds are hydrocarbons, molecules containing only carbon and hydrogen atoms. Most of the hydrocarbon molecules in gasoline are alkanes, compounds whose carbon atoms are linked together by single bonds. Most of the alkanes in gasoline contain between 6 and 11 carbon atoms. Table 4.1 lists the names and formulas of several small alkanes, along with molecular models showing their structures. The general formula for any alkane is CnH2n+2, where n is an integer. For alkanes with four or more carbon atoms, there are actually multiple possible structures having the same formula. The structures shown in Table 4.1 are known as straight-chain forms because all of the carbon atoms are linked from end to end in a single chain. Other forms involve branched chains, similar to what we saw in Section 2.8 for polyethylene. Two or more structures with the same chemical formula are called isomers. For C5H12, three possible isomers exist. In addition to the straightchain form shown in Table 4.1, we could also have the two branched forms shown below. As the number of carbons in the chain increases, the number of possible isomers increases rapidly; for C10H22, 75 isomers are possible. CH3CH2CH2CH2CH3 Pentane

The names 2-methylbutane and 2,2-dimethylpropane are examples of the systematic IUPAC nomenclature used for organic compounds. The numbers indicate the positions at which various functional groups are found.

CH3 CH3CHCH2CH3 2-Methylbutane

CH3 H3CCCH3 CH3 2,2-Dimethylpropane Structural isomers of pentane, C5H12

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Chapter 4 Stoichiometry

Table

❚

4.1

The first ten alkanes

Compound

Formula

Methane

CH4

Ethane

C2H6

Propane

C3H8

Butane

C4H10

Pentane

C5H12

Hexane

C6H14

Heptane

C7H16

Octane

C8H18

Nonane

C9H20

Decane

C10H22

Structure

4.1

Gasoline and Other Fuels

101

Gasoline actually contains several different isomers of octane and the equations we write would apply to any of those substances.

Additional products of incomplete combustion are also possible. These include formaldehyde (CH2O) and soot (particles of carbon).

Figure 4.1 ❚ This molecular scale picture of the combustion of octane to give carbon dioxide and water shows the relative number of molecules of each compound involved.

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When gasoline is burned in an engine, all of these various compounds undergo combustion simultaneously, reacting with oxygen from the air. As you might imagine, an accurate description of the combustion of such a complex mixture of compounds would be quite a challenge. So for our present purposes, it will be helpful to use a simplifying model. The simplest possible model is one in which a single compound is used to represent the gasoline mixture, and the most common choice for that compound is octane, C8H18. If we start by assuming that gasoline can be represented reasonably by octane, then it is fairly easy to write and balance a chemical equation for its combustion. If we further assume complete combustion, then our equation should show octane and oxygen as reactants, and carbon dioxide and water as products. Because we made two significant assumptions, this equation is a highly idealized model for the combustion of gasoline. The use of octane to represent all of the hydrocarbons in gasoline is mainly for simplicity. If we chose to, it would not be very difficult to write similar combustion equations for each hydrocarbon that is actually present. But the assumption of complete combustion is more drastic. You probably know that typical automobile exhaust contains a number of compounds besides carbon dioxide and water vapor. Most states require periodic emissions testing to measure the levels of carbon monoxide and hydrocarbons in a car’s exhaust, and some states or local areas also require additional tests for other types of compounds. Because these compounds don’t appear among the products in our equation above, their presence implies that our simple model does not show the full picture. What additional factors could we consider to get a more complete description of engine chemistry? We know that gasoline itself contains a wide array of hydrocarbons. So the presence of hydrocarbons in the exhaust most likely indicates that some of the gas has made it through the engine and into the exhaust stream without burning. This could mean that some hydrocarbon molecules have simply not reacted at all before leaving the engine. Or it could mean that some larger molecules have broken down into smaller fragments and that some of those fragments have made it into the exhaust. What about the presence of carbon monoxide? We wrote our original equation based on the assumption of complete combustion. This means that all carbon atoms are converted to CO2. If insufficient oxygen is present in the engine cylinders, then combustion will be incomplete, and CO is produced instead of CO2. In Figures 4.1 and 4.2 we have written two chemical equations with the same reactants but different products. How do the molecules know which reaction to 2 C8H18

⫹

25 O2

16 CO2

⫹

18 H2O

2 C8H18

⫹

17 O2

16 CO

⫹

18 H2O

Figure 4.2 ❚ Here we show the same sort of molecular picture as in Figure 4.1 but this time for the incomplete combustion of octane to give carbon monoxide and water. Note that both equations balance, but they clearly show different stoichiometry.

undergo? Under most conditions, both reactions will occur to some extent, but the relative importance of each reaction depends on the reaction conditions. The ratio of fuel to oxygen is of particular importance and is closely monitored by the fuel injection system in modern car engines. Other factors, including engine temperature, are also important. A properly tuned engine ensures that the extent of complete combustion is maximized, limits hazardous CO emissions, and also improves gas mileage. The equations we have written provide us with a good starting point in our consideration of the chemistry of gasoline or other hydrocarbon fuels. But many of the important questions surrounding the chemistry of fuels are quantitative. How much oxygen is needed to allow complete combustion to occur? How does the amount of CO2 emitted relate to the amount of fuel burned? To begin to answer questions such as these, we will need to consider the stoichiometry of the underlying chemical reactions.

4.2

Fundamentals of Stoichiometry

Stoichiometry is a term used to describe quantitative relationships in chemistry. Any chemistry question that asks “How much?” of a particular substance will be consumed or formed in a chemical reaction is a stoichiometry problem. And at the heart of every such stoichiometry problem, you’ll always find a balanced chemical equation. We have seen already that chemical equations are always written in terms of the numbers of particles involved. Whether we interpret them in terms of individual molecules or moles of molecules, the stoichiometric coefficients that balance a chemical equation refer to numbers of particles and not to masses. Usually, we can’t measure the number of particles directly in the laboratory; masses and volume of liquids are the quantities that are more likely to be measurable. Thus if we want to make quantitative calculations for a chemical reaction, frequently we need to convert between the measured value of a mass or volume and the desired value of a number of moles. Because

4.2 Fundamentals of Stoichiometry

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such calculations are common and important, chemists have developed a standard approach to overcome this mismatch of units. Although you might think of this approach as an algorithm for solving a particular class of chemistry problems, it is instructive to understand its conceptual underpinnings. The key concept is the use of the balanced chemical equation to establish ratios among moles of the various materials in the reaction.

Obtaining Ratios from a Balanced Chemical Equation The simplest combustible hydrocarbon—and therefore one of the simplest of all fuels—is methane, CH4. So the combustion of methane will provide a good place for us to begin our exploration of reaction stoichiometry. We start by writing a balanced chemical equation:

The physical state of the water formed in combustion will depend on the reaction conditions.

CH4(g)

+

2O2(g)

CO2(g)

+

2H2O(ᐉ)

In Chapter 3, we have already considered how we might read this chemical equation. We could think of the equation as representing one molecule of methane gas reacting with two molecules of oxygen or a mole of methane molecules reacting with two moles of oxygen molecules. In either case, we must have a 1:2 ratio of methane molecules to oxygen molecules, a 1:1 ratio of methane molecules to carbon dioxide molecules, a 1:2 ratio of methane molecules to water molecules, and a 2:2 (or 1:1) ratio of oxygen molecules to water molecules. As we discussed in Section 3.4, the coefficients in a chemical equation can be interpreted as molar relationships as well as molecular relationships. So we can say that the equation shows that one mole of methane reacts with two moles of oxygen to produce one mole of carbon dioxide and two moles of water. From the chemical equation, we can write the following set of mole ratios: 1 mol CH4 : 2 mol O2 1 mol CH4 : 1 mol CO2 1 mol CH4 : 2 mol H2O 2 mol O2 : 2 mol H2O We could also write the last one as 1:1. When doing stoichiometric calculations, mole ratios are very often written to look like fractions. They are used analogously to unit conversion factors, relating the amount of one substance to that of another: 1 mol CH4 2 mol O2

——————

1 mol CH4 1 mol CO2

——————

1 mol CH4 2 mol H2O

——————

2 mol O2 2 mol H2O

——————

1 mol O2 1 mol H2O

qor ——————r

Just like unit conversion factors, these mole ratios can be inverted as needed to carry out a particular calculation. Example Problem 4.1 shows how to use mole ratios in a stoichiometry problem.

E XAM PL E PRO BLEM

4 .1

In the combustion of methane, how many moles of O2 are required if 6.75 mol of CH4 is to be completely consumed?

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Strategy We start with the balanced chemical equation and use the stoichiometric coefficients to establish the mole ratio between methane and oxygen. Then, we can use that ratio to relate the amount of methane in the reaction to the amount of oxygen needed.

Solution The balanced equation was given earlier in the chapter: CH4(g) + 2 O2(g) : CO2(g) + 2 H2O(𝓵) The coefficients from this equation give us the ratio between CH4 and O2, which can be expressed in either of the following forms: 1 mol CH4 2 mol O2

2 mol O2 1 mol CH4

—————— or ——————

To calculate the desired amount of O2 from the known amount of CH4, we should use the second form. This gives us the result needed: 2 mol O2 6.75 mol CH4 × —————— = 13.5 mol O2 1 mol CH4

Discussion The answer that we obtained is a direct result of the 2:1 mole ratio and so is reasonably intuitive. You might think of the choice of the form of mole ratio as an application of dimensional analysis. Many stoichiometry problems seem more complicated than this because the known quantity is usually not in moles. But the simple use of the mole ratio illustrated here is always the pivotal step in any reaction stoichiometry calculation.

Check Your Understanding How many moles of H2O are formed in the reaction above?

The balanced chemical equation provides all the mole ratios needed to relate the amounts of the compounds in a reaction. To carry out the reaction in the laboratory, however, we often use a balance to measure out the materials needed, giving us a measurement of mass in grams. The ratios from the chemical equation are in numbers of molecules or moles. So we need to be able to convert between grams and moles. As we have seen, the molar mass provides the way to do this. The use of ratios to convert from one unit to another is no more complicated in stoichiometry than in simple physical measurements. The important difference, however, is that in many stoichiometry problems, ratios are used three or more times. Each individual use of the ratio follows principles you already know, and if you keep that in mind, you should be able to find your way through even the most complex stoichiometry problems. The main challenge is to determine which ratio to use and when to use it. One way to keep track of the steps needed is to create a flow chart using blocks to indicate when information is derived from a relationship, such as the balanced chemical equation or the molar mass. In Figure 4.3, we can see that if we begin with the mass of one species in the reaction and we want to know the mass of some other species consumed or produced, we have to use three ratios: a ratio that uses molar mass of the substance given, a mole ratio from the balanced chemical equation, and a ratio using the molar mass of the

Mass given (grams)

Molar mass RATIO

Moles in

Balanced chemical equation

Moles out RATIO

Molar mass RATIO

Mass desired (grams)

Figure 4.3 ❚ This flow diagram illustrates the various steps involved in solving a typical reaction stoichiometry problem.


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substance whose mass we wish to find. The simplest and most reliable way to identify which ratio to use for each step is to write all quantities with their complete units attached. In setting up the simple calculation in Example Problem 4.1, we wrote the initial quantity with units of “mol CH4,” and this helped us choose the correct form of the mole ratio for the calculation. Of course, keeping proper units attached to all the quantities in a calculation is always a good idea, whether it is a chemical stoichiometry problem or an engineering calculation. The following examples apply this idea to slightly more common stoichiometry calculations, where amounts are measured in mass rather than numbers of moles.

E XAM PL E PRO BLEM

4 .2

In Section 3.2, we considered the reaction between hydrogen and oxygen to form water. How many grams of water can be produced if sufficient hydrogen reacts with 26.0 g of oxygen?

Strategy First, it will help if we recognize that this is a reaction stoichiometry problem. Two signs of this are that it asks “How much . . .” and that there is obviously a chemical reaction involved. Once we see that, we should immediately realize that we need a balanced chemical equation to give us the mole ratio between water produced and oxygen reacted. We will convert the mass of oxygen given to moles of oxygen, use the mole ratio between oxygen and water to provide moles of water, and then use the molar mass of water to find the mass of water produced.

Solution Recall that both hydrogen and oxygen are diatomic gases. So the balanced equation is 2 H2(g) + O2(g) : 2 H2O(g) The molar mass of oxygen is 32.0 g/mol, and the molar mass of water is 18.0 g/mol. Using the diagram in Figure 4.3, we write sequential ratios, starting with the mass of oxygen given and keeping in mind that the answer should be in grams of water: Molar mass ratio for oxygen Mass given

Molar mass ratio for water Mole ratio from equation

26.0 g O2 ⫻ 1 mol O2 ⫻ 2 mol H2O ⫻ 18.0 g H2O ⫽ 29.3 g H2O 1 mol O2 1 mol H2O 32.0 g O2

Analyze Your Answer These two masses—for the O2 reactant and the H2O product—are similar. Does this make sense? Looking at the equation for the reaction, we can see that all of the oxygen atoms end up in water molecules. Because hydrogen is such a light element, the mass of water formed will not be too much greater than the mass of oxygen reacting. Our answer of 29.3 g of H2O is slightly larger than the 26.0 g of oxygen we started with, so it seems reasonable. (Note that this similarity between the reactant and product masses will not always be the case.)

Discussion Note that we start with the substance whose quantity was given (26.0 g O2). The first and last steps are just conversions between mass and moles, with which you should be comfortable. The use of the mole ratio from the balanced equation is the only new concept here.

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This calculation can also be carried out in a more stepwise fashion, in which you might first convert the 26.0 g O2 to 0.813 mol O2 and so forth. The result will be the same as long as all of the calculations are carried out correctly.

Check Your Understanding Calculate the mass of hydrogen needed to produce 29.2 grams of water.


4.3

Tetraphosphorus trisulfide, P4S3, is used in the manufacture of “strike anywhere” matches. Elemental phosphorus and sulfur react directly to form P4S3: 8 P4 + 3 S8 : 8 P4S3 If we have 153 g of S8 and an excess of phosphorus, what mass of P4S3 can be produced by this reaction?

Strategy The heart of any stoichiometry problem is the balanced chemical equation that provides the mole ratio we need. We must convert between masses and the number of moles in order to use this ratio, first for the reactant, S8, and at the end of the problem for the product, P4S3. Molar mass is used to provide the needed conversion factors. The phrase, “excess of phosphorus,” tells us that we have more than enough P4 to consume 153 g of S8 completely.

Solution The molar mass of S8 is 256.6 g/mol and that of P4S3 is 220.1 g/mol. Along with the balanced equation, those masses allow us to set up the needed relationships: 220.1 g P4S3 8 mol P4S3 1 mol S8 153 g S8 × ————— × —————— × —————— = 3.50 3 102 g P4S3 256.6 g S8 3 mol S8 1 mol P4S3

Check Your Understanding The smell of a burning match is due to the formation of sulfur oxides. We can represent the relevant reaction as S + O2 : SO2. If 4.8 g of sulfur is burned in this manner, what mass of sulfur dioxide is produced?

Now, we can relate the amount of carbon dioxide emitted by an automobile engine and the amount of gasoline burned. What mass of CO2 will be produced for every gallon of gasoline used? While this may sound much different from a typical chemistry book stoichiometry problem, that’s really all it is. Start by assuming that our simple model using the complete combustion of octane to represent the burning of gasoline is accurate. The balanced chemical equation is one we saw earlier: 2 C8H18 + 25 O2 : 16 CO2 + 18 H2O We’ll need to know the mass of a gallon of gasoline. As you might guess, this depends on the exact composition of the gasoline in question. But because we are basing our calculation on octane, it might be reasonable to use the density of octane, which is 0.7025 g/mL. Use that and some unit conversion factors to find the mass in grams of a gallon of octane: 0.7025 g octane 3.7854 L 1000 mL 1 gallon × ————— × ————— × ———————— = 2659 g 1L mL 1 gallon Now we’ve reduced our original question to something that looks a lot more like the previous examples: What mass of CO2 is produced by the complete combustion of 2659 g of C8H18? Use molar masses and the mole ratio from the balanced equation to set up the calculation: 44.010 g 1 mol C8H18 16 mol CO2 2659 g C8H18 × ———————— × —————— × —————— = 8196 g CO2 114.23 g C8H18 2 mol C8H18 1 mol CO2


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To put that in more familiar units, use the fact that a pound is approximately 454 g. Our calculation then says that burning a gallon of gasoline will produce 18 pounds of carbon dioxide. Although we made a few simplifying assumptions, our result should certainly be a reasonable estimate.

4.3

Limiting Reactants

When we carry out chemical reactions, the available supply of one reactant is often exhausted before the other reactants. As soon as we run out of one of the reactants, the reaction stops. The reactant completely consumed has determined how far the reaction can go, limiting the quantity of product produced. We say that the reactant completely consumed in the reaction is the limiting reactant. To illustrate this important concept, let’s imagine that somehow we can do an experiment involving just a few molecules. Suppose we start with six molecules each of H2 and O2 and allow them to react. How much water can be formed? We know the equation for this reaction from Example Problem 4.2: 2 H2(g) + O2(g) : 2 H2O(g) Figure 4.4 shows what such an experiment might look like. The reaction will proceed as long as both reactants are present. But after six molecules of water have been formed, all of the H2 will have been consumed. Even though three more molecules of oxygen remain, they cannot react. So the amount of product formed is limited by the amount of hydrogen that was available, and we say that hydrogen was the limiting reactant. In many cases, we manipulate the amounts of reactants to make certain that one compound is the limiting reactant. If the synthesis of a particular compound or material involves one scarce or expensive reactant, then a clever process engineer would most likely be sure to make that the limiting reactant by using large excesses of other reactants. Combustion reactions like burning gasoline in a car engine provide a ready example. A large excess of oxygen is always available because it can simply be drawn from the air. Your car may run out of gas, but it is not likely to run out of oxygen.

Figure 4.4 ❚ This molecular scale picture illustrates the concept of the limiting reactant. The box on the left represents the mixture of reactants in our hypothetical experiment: six H2 molecules (shown in white) and six O2 molecules (in red). The right-hand side shows the molecules present once the reaction has gone to completion: six H2O molecules and three unreacted O2 molecules. Here H2 was the limiting reactant. Because no more H2 molecules are available, no further reaction is possible.

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So in any calculation related to fuel consumption or exhaust emissions, it would be reasonable to assume that gasoline would always be the limiting reactant. In other circumstances, the identity of the limiting reactant may not be so obvious. In a laboratory experiment, we might simply mix conveniently available quantities of substances to begin to study their reaction. Determining the limiting reactant requires comparing the amount of each reactant present. As always, we should keep in mind that balanced chemical equations tell us the ratios in which moles of one substance react with moles of another. So we cannot compare masses to determine the limiting reactant; we must do the comparison using the relevant mole ratios. Example Problems 4.4 and 4.5 show how we can determine which substance is in excess and which is the limiting reactant.


4.4

A solution of hydrochloric acid contains 5.22 g of HCl. When it is allowed to react with 3.25 g of solid K2CO3, the products are KCl, CO2, and H2O. Which reactant is in excess?

Strategy As for any reaction stoichiometry problem, we should start with a balanced equation. One way to proceed from there is to calculate the amount of one reactant that would combine with the given amount of the second reactant. Comparing that with the amount actually available will reveal the limiting reactant.

There are a variety of possible strategies for identifying a limiting reactant.

Solution We know the reactants and products, so we can write a skeleton equation easily: HCl + K2CO3 : KCl + CO2 + H2O Balancing this gives us the chemical equation we need for our stoichiometry calculation: 2 HCl + K2CO3 : 2 KCl + CO2 + H2O Let’s use the quantity of HCl given and calculate the amount of K2CO3 that would react completely with it. 138.2 g K2CO3 1 mol K2CO3 1 mol HCl 5.22 g HCl × —————— × ——————— × ——————— = 9.89 g K2CO3 36.46 g HCl 2 mol HCl 1 mol K2CO3 So the given quantity of HCl (5.22 g) requires 9.89 g K2CO3, but we have only 3.25 g of K2CO3 available. The reaction will stop once all of the K2CO3 is consumed. K2CO3 is the limiting reactant, and HCl is in excess.

Analyze Your Answer It’s often hard to gain an intuitive feel for limiting reactants, so that is why it is important to learn this type of calculation. There are two factors that determine the limiting reactant—the relative molar masses and the stoichiometric mole ratio. In this case, the molar mass of K2CO3 is about three times larger than that of HCl, so it is not surprising that similar masses of the two reactants would lead to K2CO3 as the limiting reactant. Check Your Understanding Ammonia is widely used in the production of fertilizers and is also a precursor to a number of important polymers. So the formation of ammonia from its elements is one of the most important industrial chemical reactions: N2 + 3 H2 : 2 NH3 This process is usually carried out using an excess of hydrogen. If a reaction vessel contains 1.5 × 103 moles of N2, how much hydrogen would be needed to ensure that it will be in excess?

4.3 Limiting Reactants

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E XAM PL E PRO BLEM

4 .5

In Example Problem 4.3, we used the reaction that produces P4S3, one of the reactants in the combustion of a match: 8 P4 + 3 S8 : 8 P4S3 If 28.2 g of P4 is allowed to react with 18.3 g of S8, which is the limiting reactant?

Strategy We can choose either reactant and determine how much of the other reactant is required to consume it entirely. Comparing the amount calculated with the amount given in the problem will let us determine which reactant is limiting.

Solution Let’s begin with P4: 256.5 g S8 1 mol P4 3 mol S8 28.2 g P4 × ————— × ————— × ————— = 21.9 g S8 123.9 g P4 8 mol P4 1 mol S8 So, 28.2 g of P4 requires 21.9 g of S8 to react completely. We have only 18.3 g of S8, so there is not enough S8 to react with all of the P4. Therefore, S8 is the limiting reactant.

Check Your Understanding To convince yourself that you can start with either reactant, redo the example starting with the quantity of S8 given. The determination of the limiting reactant is typically just a piece of a larger puzzle. In most limiting reactant stoichiometry problems, the real goal is to determine how much product could be formed from a particular reactant mixture. Perhaps the first critical step in such a problem is to correctly recognize the fact that it is a limiting reactant situation. Whenever a particular amount of more than one reactant is specified in a stoichiometry problem, it should tell us to begin by finding the limiting reactant. Once we know which substance runs out first, we are left with an ordinary stoichiometry calculation to solve the problem. Example Problem 4.6 shows how this type of problem can be approached.

E XAM PL E PRO BLEM We will discuss the use of such additives in the closing section of this chapter.

4 .6

MTBE (methyl tert-butyl ether) has been used as an additive in gasoline. The compound is produced by reacting methanol and isobutene, according to the following equation: CH3OH + (CH3)2C"CH2 : (CH3)3COCH3 Methanol

Isobutene

MTBE

If 45.0 kg of methanol is allowed to react with 70.0 kg of isobutene, what is the maximum mass of MTBE that can be obtained?

Strategy We are given amounts of two different reactants here (methanol and isobutene). So we should realize this is a potential limiting reactant situation. We can identify the limiting reactant as in the previous examples: choose one reactant and find the amount of the other that would be needed to react with it. Comparing that result with the given amount, we can determine which of the reactants is limiting. Once we have done that, it will be straightforward to calculate the expected amount of product.

Solution We’ll need the molar masses of all three compounds. Taking care to interpret the parentheses in the molecular formulas correctly, we find the following: Methanol (CH3OH): 32.042 g/mol Isobutene ((CH3)2C"CH2): 56.107 g/mol MTBE ((CH3)3COCH3): 88.149 g/mol 110


To identify the limiting reactant, we can calculate the mass of isobutene that would react with 45.0 kg (or 45,000 g) of methanol. 1 mol methanol 56.107 g isobutene 1 mol isobutene 45,000 g methanol × ————————— × ———————— × ————————— 32.042 g methanol 1 mol isobutene 1 mol methanol = 7.88 × 104 g isobutene = 78.8 kg isobutene This tells us that to use 45.0 kg of methanol, we need at least 78.8 kg of isobutene. But we have only 70.0 kg of isobutene. So the amount of isobutene available will determine the amount of MTBE that can be formed. Isobutene is the limiting reactant and methanol is in excess. Now the problem is reduced to a simpler stoichiometry calculation starting from the mass of isobutene available. 88.149 g MTBE 1 mol isobutene 1 mol MTBE 70,000 g isobutene × ————————— × ———————— × ———————— 56.107 g isobutene 1 mol isobutene 1 mol MTBE = 1.10 × 105 g MTBE = 1.10 3 102 kg MTBE

Analyze Your Answer This value is larger than the original amount of either reactant. Does this make sense? If we look closely at the reaction, it is a synthesis where the two molecules become one. If our answer had been greater than the sum of the two reactant masses, we would have cause for concern. But this value is certainly plausible.

Discussion This problem combines two types of calculations that we’ve seen in earlier examples. The first step is identifying the limiting reactant, and the second is the actual stoichiometry calculation. Note that here isobutene was the limiting reactant even though the available mass of isobutene was greater than that of methanol. This is a reminder that reaction stoichiometry calculations should always be handled in terms of moles, not masses.

Check Your Understanding The compound diborane, B2H6, was once proposed for use as a rocket propellant. It can be produced by the following reaction 3 LiBH4 + BF3 : 2 B2H6 + 3 LiF If 24.6 g of LiBH 4 is combined with 62.4 g of BF 3, what mass of diborane can be formed?

The concept of limiting reactants has some important implications for engineering applications where chemical reactions are involved. We’ve already mentioned that a good design for a chemical process would ensure that a scarce or expensive substance would be the limiting reactant, so that none of it would go to waste. A rather different situation exists for the design of rocket engines. Here the total mass of the rocket is a major consideration, so the aim would likely be to use the minimum mass of fuel that could deliver the needed thrust. Many rocket engines rely on so-called two-component fuel mixtures, where energy is released as two compounds are allowed to react. The optimal design for such a rocket would generally ensure that the two compounds are present in a stoichiometric mixture, so that neither is in excess. The following example illustrates how such a mixture might be chosen.


4.7

The solid fuel booster rockets of the space shuttle are based on the following reaction between ammonium perchlorate and aluminum: 3 NH4ClO4(s) + 3 Al(s) : Al2O3(s) + AlCl3(g) + 3 NO(g) + 6 H2O(g) 4.3 Limiting Reactants

111

If either reactant is in excess, unnecessary mass will be added to the shuttle, so a stoichiometric mixture is desired. What mass of each reactant should be used for every kilogram of the fuel mixture?

Strategy We want to ensure that there will be no excess of either reactant. Because we are asked for the composition “for every kilogram” of fuel, we might start by assuming the total mass of fuel will be 1000 g. We can write that as an equation with two unknowns: mNH ClO + mAl = 1000 g 4

4

Because we have two unknowns, we will need a second relationship between the two masses. The balanced chemical equation shows us that the mole ratio between the two reactants should be 3:3, or 1:1. We can use molar masses to convert that to a ratio in terms of masses and then use that ratio as a second equation to solve the problem.

Solution The molar mass of NH4ClO4 is 117.49 g/mol, and that of Al is 26.98 g/mol. To ensure that we have a stoichiometric mixture, we need equal molar amounts of each reactant. So if we were to use 117.49 g (1 mole) of NH4ClO4, we would need 26.98 g (also 1 mole) of Al. We can write this as an equation: mNH ClO

117.49 g 26.98 g

————— = ————— = 4.355 m 4

4

Al

And although our actual amounts will need to be larger than those, that ratio of masses will still be right. So now we have two equations for the two unknown masses, and we can solve by doing a little bit of algebra. Our two equations are mNH ClO + mAl = 1000 g 4

4

mNH ClO = 4.355 mAl 4

4

Substituting the right-hand side of the second equation in the first equation gives: 4.355 mAl + mAl = 1000 g Simplifying that and solving, we get the mass of aluminum needed for each kilogram of the fuel mixture. 5.355 mAl = 1000 g 1000 g mAl = ———— = 186.8 g Al 5.355 Now the mass of NH4ClO4 is easy to obtain. mNH ClO = 1000 g − mAl 4

4

= 1000 g − 186.8 g = 813.2 g So each kilogram of fuel should comprise 186.8 g Al and 813.2 g NH4ClO4.

Discussion The mass of fuel required for a space shuttle engine would clearly be much larger than a kilogram, but this result could easily be scaled up to the needed mass. As for most problems, there are a variety of other approaches that could be used to solve this. The particular strategy that we’ve used is based on the general idea that if we have two unknown quantities (like the two masses here), we need to find two separate relationships between them to solve the problem.

Check Your Understanding A fuel mixture comprising 600.0 g of NH4ClO4 and 400.0 g of Al is used in a test of a laboratory scale mock-up of a shuttle engine. When the engine stops burning, some unused fuel remains. What substance is this unburned fuel, and what mass of it should be found?

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4.4

Theoretical and Percentage Yields

For chemical reactions and especially for reactions that produce commercial products we would generally prefer to have reactions that are efficient, producing as much of the desired products as possible while minimizing any unwanted byproducts. Why is this important? In many industrial situations, one big reason is economics. If a business can increase the amount of product obtained from a given amount of reactants, it should become more profitable. Many factors such as temperature of the reaction, the possibility of side reactions, or further reaction of the product to form something else can decrease the amount of desired product obtained. Minimizing these unwanted reactions could also reduce the amount of waste produced, thereby lowering environmental impacts and associated costs. We can rate the efficiency of a reaction by calculating how much product would form under perfect or ideal conditions and then comparing the actual measured result with this ideal. The ideal amount of product is called the theoretical yield, and it is obtained by working a stoichiometry problem. Measuring the amount of product formed gives us the actual yield. From the ratio of the actual yield to the theoretical yield, we can calculate the percentage yield. actual yield Percentage yield = q————————r × 100% theoretical yield

The principles of green chemistry, which we introduced in Section 3.6, are increasingly important in the design of industrial chemical reactions.

(4.1)

Increasing this percentage yield is one of the foremost goals for a chemical engineer working on an industrial process.


4.8

The Solvay process is important in the commercial production of sodium carbonate (Na2CO3), which is used in the manufacture of most glass. The last step in the Solvay process is the conversion of NaHCO3 (sodium bicarbonate, or baking soda) to Na2CO3 by heating. 2 NaHCO3(s)

Na2CO3(s) + CO2(g) + H2O(g)

In a laboratory experiment, a student heats 42.0 g of NaHCO3 and determines that 22.3 g of Na2CO3 is formed. What is the percentage yield of this reaction?

Strategy We know the actual yield from the experiment. To calculate the percentage yield, first we need to find the theoretical yield. We can do that by calculating the maximum quantity of product that could form, based on the stoichiometry of the reaction. Once we have both the theoretical yield and the actual yield, finding the percentage yield is straightforward.

Solution Begin by working the stoichiometry problem. 106.0 g Na2CO3 1 mol NaHCO3 1 mol Na2CO3 42.0 g × ———————— × ———————— × ———————— = 26.5 g Na2CO3 84.0 g NaHCO3 2 mol NaHCO3 1 mol Na2CO3 Then compute the percentage yield. actual yield Percentage yield = ———————— × 100% theoretical yield 22.3 g = ———— × 100% 26.5 g = 84.2%

Analyze Your Answer A quick estimate of the answer can tell us that we have done the stoichiometry correctly. Because 42 is half of 84, we have 0.5 mol of NaHCO3.

4.4 Theoretical and Percentage Yields

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From the equation, the mole ratio of NaHCO3 to Na2CO3 is 2:1, so we would expect to obtain 0.25 mol of Na2CO3. One-fourth of 106 is about 26, so we know that the stoichiometry is correct. Because the actual yield is less than the theoretical yield, an answer less than 100% is plausible. (We may sometimes see yields greater than 100% for reactions like this, but that indicates that the product is not completely dry or contains some impurity.)

Discussion We might ask whether or not this is a “good” percentage yield. It isn’t possible to set an absolute numerical target for percentage yields. If a reaction leads to a highly desired (i.e., expensive) product for which there is no alternative source, then a yield of just a few percent might be considered adequate. In the present case, the reaction is rather simple and the product is relatively inexpensive. So a higher percentage yield would probably be needed for an economically viable plant.

Check Your Understanding In the “Check Your Understanding” of Example Problem 4.6, you considered a reaction to produce diborane. An alternative route to this propellant is shown below: 3 NaBH4 + 4 BF3 : 2 B2H6 + 3 NaBF4 If 173.2 g of BF 3 reacts with excess NaBH 4 to give 28.6 g of B 2H 6, what is the percentage yield?

How does the idea of percentage yield apply to the chemistry of fuels? For a hydrocarbon such as gasoline, complete combustion liberates more energy than incomplete combustion or other side reactions. Complete combustion also reduces the levels of several potentially harmful compounds, including CO and soot. So we might define the efficiency in terms of the extent of complete combustion. If we wanted to monitor this experimentally, one option would be to monitor the amount of CO2 in the exhaust stream. Similarly, the presence of CO, CH2O, or other carbon-containing compounds in the exhaust would allow us to monitor the extent of various side reactions. Such measurements are routinely carried out on test engines to determine the optimum operating conditions.

4.5


In the stoichiometry problems we have discussed so far, the numbers of moles involved have been obtained by using the ratio established by the molar mass of the substances. Now we can expand the range of problems that we can consider by using Equation 3.2 to determine the number of moles of substances when volumes are measured rather than masses. The heart of the stoichiometry problem, however, remains the same. The balanced chemical equation still provides the critical ratios among numbers of moles of various species in the reaction. Again, a flow diagram for solving a stoichiometry problem shows the manipulations we need to consider (Figure 4.5).

Volume and concentration (given)

n=M⫻V

Moles in

Balanced chemical RATIO equation

Moles out

n=M⫻V

Desired answer (molarity or volume)

Figure 4.5 ❚ This flow diagram shows the important steps in a typical solution stoichiometry calculation.

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The only complication added by considering reactions in solution is the need to relate three variables, using Equation 3.2, rather than two variables using the molar mass ratio. Example Problem 4.9 shows how to approach these solution stoichiometry problems.


4.9

As we mentioned in Example Problem 3.10, the fuel hydrazine can be produced by the reaction of solutions of sodium hypochlorite and ammonia. The relevant chemical equation is NaClO(aq) + 2 NH3(aq) : N2H4(aq) + NaCl(aq) + H2O(𝓵) If 750.0 mL of 0.806 M NaClO is mixed with excess ammonia, how many moles of hydrazine can be formed? If the final volume of the resulting solution is 1.25 L, what will be the molarity of hydrazine?

Strategy We are asked for the expected amount of a product, so this is a reaction stoichiometry problem. Because NH3 is said to be in excess, we know that NaClO will be the limiting reactant. So we will use the given volume and concentration to find the number of moles of NaClO reacting. Then use the mole ratio from the balanced equation to find the number of moles of N2H4 that can be formed. Finally, we can use that number of moles and the given final volume to obtain the molarity.

Solution Convert 750.0 mL to 0.7500 L and determine the number of moles of NaClO reacting: nNaClO = M × V = 0.806 mol/L × 0.7500 L = 0.605 mol NaClO Use that along with the 1:1 mole ratio from the balanced equation to find the corresponding number of moles of hydrazine: 1 mol N2H4 0.605 mol NaClO × ——————— = 0.605 mol N2H4 1 mol NaClO To answer the second part of the problem, we just need to use this number of moles along with the given final volume to find the molarity. 0.605 mol N2H2 n = ———————— M = —— = 0.484 M N2H4 1.25 L V

Discussion Notice that the stoichiometric calculation is done using the mole ratio from the balanced equation, just as in all of our earlier examples. Only the conversions to and from numbers of moles have changed because this reaction is in solution.

Check Your Understanding What volume (in milliliters) of 0.150 M HCl is required to react completely with 0.503 g of dry Na2CO3? A common laboratory technique, called titration, requires understanding solution stoichiometry. A solution-phase reaction is carried out under controlled conditions so that the amount of one reactant can be determined with high precision. A carefully measured quantity of one reactant is placed in a beaker or flask. A dye called an indicator can be added to the solution. The second reactant is added in a controlled fashion, typically by using a burette (Figure 4.6). When the reaction is complete, the indicator changes color. When the indicator first changes color, we have a stoichiometric mixture of reactants. We know the number of moles of the first reactant (or the molarity and volume) and the volume of the second reactant used. So as long as we know the balanced equation for the reaction, we can find the unknown concentration of the second reactant.

Indicators are also examples of weak acids or bases, as defined in Chapter 3.

4.5


115


Figure 4.6 ❚ The photo sequence shows some of the steps in a typical titration. One solution is poured into a burette in the left-hand panel. The second solution is placed into an Erlenmeyer flask, and the burette is positioned above the flask in the second panel. The valve on the burette allows controlled addition of solution, and the shape of the Erlenmeyer flask permits easy mixing by swirling. In the final panel, the solution in the flask turns pink when the endpoint is reached.

E XAM PL E PRO BLEM

4 .10

Many common titrations involve the reaction of an acid with a base. If 24.75 mL of 0.503 M NaOH solution is used to titrate a 15.00-mL sample of sulfuric acid, H2SO4, what is the concentration of the acid?

Strategy A titration problem is an applied stoichiometry problem, so we will need a balanced chemical equation. We know the molarity and volume for the NaOH solution, so we can find the number of moles reacting. The mole ratio from the balanced equation lets us calculate moles of H2SO4 from moles of NaOH. Because we know the volume of the original H2SO4 solution, we can find its molarity.

Solution The acid–base reaction will produce sodium sulfate and water. The balanced equation is H2SO4(aq) + 2 NaOH(aq) : Na2SO4(aq) + 2 H2O(𝓵) Determine the number of moles of NaOH from the given molarity and volume: 0.503 mol NaOH 0.02475 L solution × ————————— = 0.0124 mol NaOH 1 L solution Use the balanced chemical equation to determine the number of moles of H2SO4: 1 mol H2SO4 0.0124 mol NaOH × ——————— = 6.22 × 10−3 mol H2SO4 2 mol NaOH Determine the concentration of H2SO4: 6.22 × 10−3 mol H2SO4 M = ——————————— = 0.415 M H2SO4 0.01500 L solution

Analyze Your Answer The molarity we found for sulfuric acid is smaller than that we had for sodium hydroxide. Does this make sense? The point to remember here is the stoichiometry that provides the 2:1 mole ratio. Even though the volume of NaOH used is greater than the volume of the acid, it is not twice the volume, so our H2SO4 molarity should be smaller.

Discussion This is the type of problem that students get wrong if they use the dilution formula (Eq. 3.3) by mistake. The dilution formula does not include the mole ratio from a balanced chemical equation, and in this case, the 2:1 ratio would result in an error of a factor of 2 if you tried to use Eq. 3.3.

Check Your Understanding Analysis of a solution of NaOH showed that 42.67 mL of 0.485 M HNO3 was needed to titrate a 25.00-mL sample of NaOH. What is the concentration of the NaOH solution?

116


INSIGHT INTO

4.6

Alternative Fuels and Fuel Additives

In the opening insight, we pointed out that gasoline is a rather complicated mixture of chemical compounds. Most of these are hydrocarbons of one sort or another, whose origins can be traced back to the oil from which the gasoline was refined. But the gasoline you buy is also likely to contain a range of fuel additives. Specific reasons for using these fuel additives generally fall under three broad categories: improving engine performance, reducing undesirable engine emissions, or reducing dependence on imported petroleum products. Some additives can achieve more than one of these goals simultaneously. Let’s first examine how additives can improve performance. In the gasoline industry, different grades of fuel are usually described in terms of their octane ratings; higher octane numbers are expected to deliver better engine performance. The octane number is really a measure of how highly the gas can be compressed in the engine cylinder before it will spontaneously ignite. This allows an engine to operate at higher compression without unwanted “knocking,” which occurs if the fuel mixture in a cylinder ignites prematurely. Some carmakers recommend that only premium gas be used in their cars because the engines are designed to run at higher than usual compression ratios. The term “octane number” has its origins in the fact that octane itself is highly compressible; one way to achieve a higher octane number is to increase the fraction of octane in a particular gasoline. But what is important is the compressibility of the gas, not the amount of octane present. There are a number of ways to improve compressibility that are easier and cheaper than boosting the actual octane content. The first widely used fuel additive was tetraethyl lead, which substantially increases compressibility. From the 1920s into the 1970s, practically all gasoline sold in the United States contained lead additives. But the increased performance and reduced fuel costs came at a price. Lead is toxic, and most of the lead introduced into gasoline was ultimately emitted into the atmosphere and settled into soil and water. Although these health concerns were known early, they did not receive widespread attention until heightened environmental interest took hold in the 1970s. The final blow to leaded gas came with the introduction of the catalytic converter in the 1970s. As you may know, the catalytic converter is designed to reduce the emission of carbon monoxide and other pollutants. The important chemistry of the converter takes place on a platinum metal surface. But the microscopic lead particles in the exhaust stream “poison” the platinum surface by binding to it so strongly that the catalytic reaction cannot take place. So “unleaded” gas was introduced and gradually took over the market; since 1996 it has been illegal to sell gasoline containing lead additives for use as a motor fuel in the United States. Other additives have stepped in to take on the octane-boosting role formerly played by tetraethyl lead. The most widely used in the United States today is ethanol. Because adding ethanol increases the oxygen content of gasoline, it is often referred to as an oxygenate, and gasoline containing ethanol is known as oxygenated fuel. The added oxygen helps to ensure more complete combustion and therefore reduces the emission of carbon monoxide, hydrocarbons, and soot. Gasoline containing at least 2% oxygen by weight is known as reformulated gasoline (RFG), and the use of such fuels is mandatory in some areas where pollution problems are severe. Ethanol for use as a fuel additive is produced from crops such as corn, barley, and wheat. Starch from these crops can be converted easily to sugar, and the sugar is then fermented to produce ethanol. Because the feedstocks are agricultural, ethanol production does not rely on imported petroleum products. Gasoline containing up to 10% ethanol can be burned in any modern automobile, and mixtures containing up to 85% ethanol can be used in specially designed engines. When used at such high levels, ethanol would more properly be thought of as an alternative fuel rather than a gasoline additive. Ethanol use is most common in the midwestern United States. Another additive, MTBE (methyl tert-butyl ether, (CH3)3COCH3), was widely used in the late 1990s. (See Example Problem 4.6.) Like ethanol, MTBE increases

We will discuss catalysis and catalytic converters in Chapter 11.

Gasoline formulations are typically seasonal. This helps explain the commonly observed increase in gasoline prices in late spring when refineries are changing from winter formulations to summer formulations and demand outpaces supply during the transition.

4.6 Alternative Fuels and Fuel Additives

117

Figure 4.7 ❚ Structural formulas of tetraethyl lead and MTBE.

CH3 CH2 H3C

CH2 Pb CH2 CH3 CH2 CH3 Tetraethyl lead

A decision on the introduction of a new compound like MTBE into widespread use is always made on the basis of incomplete information. Engineers and scientists must rely on the best data available when the decision is being made.

CH3 H3C

C

CH3

O CH3 MTBE

both the octane rating and oxygen content of gasoline, and fuel with as much as 15% MTBE can be used in any modern engine. After the passage of the 1990 Clean Air Act, the use of MTBE increased dramatically. The law mandated the use of oxygenated fuels, and the fuel industry saw that MTBE offered advantages over ethanol in cost of production and transportation. So the information available at the time suggested that introducing MTBE made sense in both environmental and business terms. Later, however, possible health concerns over MTBE called that choice into doubt because MTBE is a possible carcinogen, meaning it may cause some forms of cancer. As the usage of MTBE in gasoline increased, reports of detectable levels of the compound in groundwater also became more widespread. The source of this MTBE was generally believed to be leaks from pipelines or storage tanks. (Unlike lead, MTBE and ethanol undergo combustion along with the other components of gasoline, so no significant levels of these additives are found in vehicle exhaust streams.) Although the extent of MTBE’s impact on human health is still somewhat uncertain, some areas (including the state of California) have banned the use of MTBE in gasoline.

FOCU S O N

PRO BLEM S O LVI NG

Question In the 1950s, both the United States and the Soviet Union had programs to use decaborane, B10H14, as a rocket fuel. This chemical reacts violently with oxygen according to the reaction, B10H14 + 11 O2 : 5 B2O3 + 7 H2O. For maximum efficiency in rocket design, an engineer devises a design that allows both fuels (in this case, decaborane and liquid oxygen) to run out at roughly the same time. Suppose that a particular rocket design allows for a total of 1.20 × 106 kg of the two fuels. Describe how you would determine how much fuel should be decaborane and how much should be oxygen.

Strategy This is a type of stoichiometry problem because the two components of the fuel react according to the chemical equation that is provided. There are several ways that this problem could be approached, but all would have to deal with the fact that the chemical equation relates moles of materials, and the problem must be addressed in terms of masses of materials. We must devise a solution with two equations in two unknowns to obtain this answer.

Solution The first equation should take advantage of the fact that there are only two reactants. Then, the desired total mass must be the sum of the masses of the two components: Mass B10H14 + mass O2 = 1.20 × 106 kg The second equation can be determined from the reaction mole ratios and the molar masses of the chemicals involved. (Those could be looked up and calculated as usual.) 1 mol B10H18 molar mass B10H18 1 mol O2 Mass B10H14 = mass O2 × ——————— × ——————— × ————————— molar mass O2 11 mol O2 1 mol B10H18 Assuming we have values for the needed molar masses, we can solve these two equations in two unknowns to obtain the answer.

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S U M M ARY The quantitative relationships associated with chemical reactions are vitally important for many applications of chemistry— including engineering of internal combustion engines and the fuels that power them. A balanced chemical equation is the key to any quantitative understanding of a reaction. Quantitative calculations for chemical reactions are referred to as stoichiometry problems, and the most important idea in this chapter is that the heart of every such problem is the balanced chemical equation. The balanced chemical equation provides mole ratios for all reactants and products. To use these ratios, we usually have to convert information from something that can be easily measured—such as the mass or volume of a reactant or product—to moles. This establishes the main pattern for solving stoichiometry problems. First, we must look at the information we know about a reactant or product and convert that to moles of that substance. We use the balanced chemical equation and the mole ratios it contains to convert from the known molar

amount of one substance to the number of moles of the substance we wish to know about. Finally we usually convert from moles to mass (or volume) of the unknown substance. All stoichiometry problems can be approached with this general pattern. In some cases, however, additional calculations may be needed. For example, if we are given (or able to measure) known amounts of two or more reactants, we must determine which of them will be completely consumed (the limiting reactant). Once again, the mole ratios in the balanced equation hold the key. Another type of calculation that can be considered in a stoichiometry problem is determining the percentage yield of a reaction. In this case, the amount of product determined in the problem represents the theoretical yield of the reaction—a value that is not often attained when a reaction is carried out in the laboratory or in industry. If the actual yield can be measured separately, the ratio of the actual yield to the theoretical yield provides the percentage yield.

K EY TE RM S actual yield (4.4)

limiting reactant (4.3)

reformulated gasoline (RFG) (4.6)

alkanes (4.1)

side reactions (4.4)

fuel additives (4.6)

methyl tert-butyl ether (MTBE) (4.6)

hydrocarbons (4.1)

mole ratio (4.2)

theoretical yield (4.4)

indicator (4.5)

oxygenated fuels (4.6)

titration (4.5)

isomers (4.1)

percentage yield (4.4)

stoichiometry



(a) 2 H2 + O2 : 2 H2O (b) 2 H2O2 : 2 H2O + O2

INSIGHT INTO Gasoline and Other Fuels 4.1 List at least two factors that make it difficult to describe the combustion of gasoline accurately. What assumptions can be made to address these complications? 4.2 What is an alkane? 4.3 Explain the difference between complete and incomplete combustion. 4.4 Automobile exhaust often contains traces of formaldehyde (CH2O), which is another possible product of incomplete combustion. Write a balanced equation for the formation of formaldehyde during the combustion of octane. (Water will also be formed as a product.) 4.5 Methane, ethane, and propane are also hydrocarbons, but they are not major components of gasoline. What prevents them from being part of this mixture? 4.6 Use the web to research prices of gasoline at the pump for consumers in your area during the past year. Also find information about the price of crude oil for the same period. Discuss any correlations you observe between the two prices. Fundamentals of Stoichiometry 4.7 For the following reactions, write the ratios that can be established among molar amounts of the various compounds:

■


(c) P4 + 5 O2 : P4O10 (d) 2 KClO3 : 2 KCl + 3 O2 4.8 In an experiment carried out at very low pressure, 1.3 × 1015 molecules of H2 are reacted with acetylene, C2H2, to form ethane, C2H6, on the surface of a catalyst. Write a balanced chemical equation for this reaction. How many molecules of acetylene are consumed? 4.9

4.10

Sulfur, S 8 , combines with oxygen at elevated temperatures to form sulfur dioxide. (a) Write a balanced chemical equation for this reaction. (b) If 200 oxygen molecules are used up in this reaction, how many sulfur molecules react? (c) How many sulfur dioxide molecules are formed in part (b)?

■

■ How many moles of oxygen can be obtained by the decomposition of 7.5 mol of reactant in each of the following reactions?

(a) 2 KClO3 : 2 KCl + 3 O2 (b) 2 H2O2 : 2 H2O + O2 (c) 2 HgO : 2 Hg + O2 (d) 2 NaNO3 : 2 NaNO2 + O2 (e) KClO4 : KCl + 2 O2


119

4.11 MTBE, C 5H 12O, is one of the additives that replaced tetraethyl-lead in gasoline. (See Example Problem 4.6 and Section 4.6.) How many moles of O2 are needed for the complete combustion of 1.50 mol of MTBE? 4.12 In petroleum refining, hydrocarbons are often manipulated by reacting them with H2(g). If hexene, C6H12, is reacted with hydrogen to form hexane, C6H14, how many moles of hydrogen are needed to react with 453 moles of hexene? 4.13

■ For

the following reactions, determine the value of x.

(a) 4 C + S8 : 4 CS2 3.2 mol S8 yield x mol CS2 (b) CS2 + 3 O2 : CO2 + 2 SO2 1.8 mol CS2 yield x mol SO2 (c) N2H4 + 3 O2 : 2 NO2 + 2 H2O 7.3 mol O2 yield x mol NO2 (d) SiH4 + 2 O2 : SiO2 + 2 H2O 1.3 × 10−3 mol SiH4 yield x mol H2O 4.14

■ The combustion of liquid chloroethylene, C2H3Cl, yields carbon dioxide, steam, and hydrogen chloride gas. (a) Write a balanced equation for the reaction. (b) How many moles of oxygen are required to react with 35.00 g of chloroethylene? (c) If 25.00 g of chloroethylene reacts with an excess of oxygen, how many grams of each product are formed?

4.15

■ What mass of the unknown compound is formed in the following reactions? (Assume that reactants for which amounts are not given are present in excess.)

(a) C2H4 + H2 : C2H6 3.4 g C2H4 reacts to produce x g C2H6 (b) CS2 + 3 Cl2 : CCl4 + S2Cl2 5.78 g Cl2 reacts to produce x g S2Cl2 (c) PCl3 + 3 H2O : H3PO3 + 3 HCl 3.1 mg PCl3 reacts to produce x mg HCl (d) B2H6 + 3 O2 : B2O3 + 3 H2O 4.6 kg B2H6 reacts to produce x kg B2O3 4.16

4.20 Assuming a charcoal briquette is composed entirely of carbon, what mass of oxygen is needed to completely burn a 37.3-g briquette? 4.21 Ammonium nitrate, NH4NO3, will decompose explosively to form N2, O2, and H2O, a fact that has been exploited in terrorist bombings. What mass of nitrogen is formed by the decomposition of 2.6 kg of ammonium nitrate? Limiting Reactants 4.22 If 3.4 mol Al and 6.2 mol Fe2O3 are mixed, what is the limiting reactant? 2 Al + Fe2O3 : Al2O3 + 2 Fe 4.23 If 8.4 moles of disilane, Si2H6, is combined with 15.1 moles of O2, which is the limiting reactant? 2 Si2H6 + 7 O2 : 4 SiO2 + 6 H2O 4.24 Generally, an excess of O2 is needed for the reaction Sn + O2 : SnO2. What is the minimum number of moles of oxygen required to oxidize 7.3 moles of tin? 4.25 In the reaction of arsenic with bromine, AsBr5 will form only when excess bromine is present. Write a balanced chemical equation for this reaction. Determine the minimum number of moles of bromine that are needed if 9.6 moles of arsenic is present. 4.26

If 112 g CaO reacts with 224 g NH4Cl, how many moles of reactants and products are there when the reaction is complete? 4.27 When octane is combusted with inadequate oxygen, carbon monoxide may form. If 100 g of octane is burned in 200 g of O2, are conditions conducive to forming carbon monoxide? 4.28

Many metals react with halogens to give metal halides. For example, iron reacts with chlorine to give iron(II) chloride, FeCl2.

4.18 A gas is either CH4 or C2H6. A 1.00-g sample of this gas produces 1.80 g of water when combusted in excess O2. Assuming complete combustion to CO2 and H2O, which gas is it? 4.19 How many metric tons of carbon are required to reduce 7.83 metric tons of Fe2O3 according to the following reaction? 2 Fe2O3 + 3 C : 3 CO2 + 4 Fe How many metric tons of iron are produced?

120


The equation for one of the reactions in the process of turning iron ore into the metal is ■

Fe2O3(s) + 3 CO(g) : 2 Fe(s) + 3 CO2(g) If you start with 2.00 kg of each reactant, what is the maximum mass of iron you can produce?

Fe(s) + Cl2(g) : FeCl2(s)

4.17 Phosgene is a highly toxic gas that has been used as a chemical weapon at times in the past. It is now used in the manufacture of polycarbonates, which are used to make compact discs and plastic eyeglass lenses. Phosgene is produced by the reaction, CO + Cl2 : COCl2. Given an excess of carbon monoxide, what mass of chlorine gas must be reacted to form 4.5 g of phosgene?

gas can be prepared by the reaction

CaO(s) + 2 NH4Cl(s) : 2 NH3(g) + H2O(g) + CaCl2(s)

■

Beginning with 10.0 g iron, what mass of Cl2, in grams, is required for complete reaction? What quantity of FeCl2, in moles and in grams, is expected?

■ Ammonia

4.29 Copper reacts with sulfuric acid according to the following equation: 2 H2SO4 + Cu : CuSO4 + 2 H2O + SO2 How many grams of sulfur dioxide are created by this reaction if 14.2 g of copper reacts with 18.0 g of sulfuric acid? 4.30

■ One of the steps in the manufacture of nitric acid is the oxidation of ammonia shown in this equation:

4 NH3(g) + 5 O2(g) : 4 NO(g) + 6 H2O(g) If 43.0 kg of NH3 reacts with 35.4 kg of O2, what mass of NO forms? 4.31 When Al(OH)3 reacts with sulfuric acid, the following reaction occurs: 2 Al(OH)3 + 3 H2SO4 : Al2(SO4)3 + 6 H2O If 1.7 × 10 3 g of Al(OH) 3 is combined with 680 g of H2SO4, how much aluminum sulfate can form?

■


4.32

■ Copper

reacts with nitric acid via the following equation:

3 Cu(s) + 8 HNO3(aq) : 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(𝓵) What mass of NO(g) can be formed when 10.0 g of Cu reacts with 115 g HNO3?

How many grams of sulfur would be needed to obtain 80.0 g of CS2? 4.44 Sulfur hexafluoride is a very stable gas useful in electrical generators and switches. It is formed by direct synthesis from the elements. S(s) + 3 F2(g) : SF6(g)

4.33 How much HNO3 can be formed in the following reaction if 3.6 kg of NO2 gas is bubbled through 2.5 kg of water?

If 92 g of SF6 is produced from the reaction of 115 g of sulfur in excess fluorine, what is the percentage yield?

3 NO2(g) + H2O(𝓵) : 2 HNO3(aq) + NO(g) 4.34 Hydrogen and oxygen are reacted and the water formed is collected at 25°C, where it has a density of 0.997 g/mL. If 36.8 g of H2 and 168 g of O2 are reacted, how many mL of water will be collected? 4.35

■ Silicon carbide, an abrasive, is made by the reaction of silicon dioxide with graphite (solid carbon).

SiO2 + C

SiC + CO (balanced?)

We mix 150.0 g of SiO2 and 101.5 g of C. If the reaction proceeds as far as possible, which reactant is left over? How much of this reactant remains?

Sulfur hexafluoride, SF6

4.45 Magnesium nitride forms in a side reaction when magnesium metal burns in air. This reaction may also be carried out in pure nitrogen.

Theoretical and Percentage Yields 4.36 Define the term theoretical yield.

3 Mg(s) + N2(g) : Mg3N2(s)

4.37 What types of events happen in chemical reactions that lead to percentage yields of less than 100%? 4.38 Sometimes students in chemistry labs determine percentage yields greater than 100%. Assuming there is no calculation error, how could this happen? 4.39 The theoretical yield and the actual yield for various reactions are given below. Determine the corresponding percentage yields.

Reaction 1

Theoretical Yield

Actual Yield

35.0 g

12.8 g

Reaction 2

9.3 g

120 mg

Reaction 3

3.7 metric tons

1250 kg

Reaction 4

40.0 g

41.0 g

4.40 A reaction that produced 4.8 mg of taxol, an anticancer drug, is reported to have a yield of 38%. What was the theoretical yield? 4.41

4.46 Industrial production of hydrogen gas uses the reaction shown below. If 1.00 metric ton of propane reacting with excess water yields 270 kg of H2, what is the percentage yield? C3H8(g) + 3 H2O(𝓵) : 3 CO(g) + 7 H2(g) 4.47 If 21 g of H2S is mixed with 38 g of O2 and 31 g of SO2 form, what is the percentage yield? 2 H2S + 3 O2 : 2 SO2 + 2 H2O 4.48 A mixture of 10.0 g of NO and 14.0 g of NO2 results in the production of 8.52 g of N2O3. What is the percentage yield? NO(g) + NO2(g) : N2O3(𝓵) Solution Stoichiometry 4.49 Small quantities of hydrogen gas can be prepared by the following reaction: Zn(s) + H2SO4(aq) : ZnSO4(aq) + H2(g)

■ Methanol,

CH3OH, is used in racing cars because it is a clean-burning fuel. It can be made by this reaction: CO(g) + 2 H2(g) : CH3OH(𝓵)

What is the percent yield if 5.0 × 103 g H2 reacts with excess CO to form 3.5 × 103 g CH3OH? 4.42

If 18.4 g of Mg3N2 forms from the reaction of 20.0 g of magnesium with excess nitrogen, what is the percentage yield?

■ When

iron and steam react at high temperatures, the following reaction takes place. 3 Fe(s) + 4 H2O(g) : Fe3O4(s) + 4 H2(g)

How many grams of H2 can be prepared from 25.0 mL of 6.00 M H2SO4 and excess zinc? 4.50 Aluminum hydroxide can be dissolved in hydrochloric acid to form aluminum chloride and water. Write a balanced chemical equation for this reaction. How many mL of 3.37 M HCl is needed to react with 10.0 g of Al(OH)3? 4.51 What is the role of an indicator in a titration? 4.52

■ What volume of 0.812 M HCl, in milliliters, is required to titrate 1.45 g of NaOH to the equivalence point?

4.53

What volume, in milliliters, of 0.512 M NaOH is required to react completely with 25.0 mL 0.234 M H2SO4?

How much iron must react with excess steam to form 897 g of Fe3O4 if the reaction yield is 69%? 4.43

■ The

percent yield of the following reaction is consistently 87%. CH4(g) + 4 S(g) : CS2(g) + 2 H2S(g)

■


NaOH(aq) + HCl(aq) : H2O(𝓵) + NaCl(aq) ■


121

4.54

■ What

is the molarity of a solution of nitric acid if 0.216 g of barium hydroxide is required to neutralize 20.00 mL of nitric acid?

4.55

■ Hydrazine,

N2H4, a base like ammonia, can react with an acid such as sulfuric acid. 2 N2H4(aq) + H2SO4(aq) : 2 N2H5+(aq) + SO4–2(aq)

What mass of hydrazine can react with 250. mL 0.225 M H2SO4? INSIGHT INTO Alternative Fuels and Fuel Additives 4.56 Why are fuel additives used? 4.57 What is actually measured by the octane ratings of different grades of gasoline? 4.58 Offer an explanation as to how oxygenated fuels might increase the extent of combustion. 4.59 Perform a web search to learn what restrictions (if any) on the use of MTBE in gasoline might exist in your local area. 4.60 Using the web, find information about the amount of lead in the environment during the past 40 years. Correlate what you observe with the presence or absence of tetraethyl lead in gasoline. 4.61 Using the web, find out how lead “poisons” the catalyst in a catalytic converter. Additional Problems 4.62 Estimate the mass of CO2 emitted by a typical automobile in a year. 4.63

■ You

have 0.954 g of an unknown acid, H2A, which reacts with NaOH according to the balanced equation H2A(aq) + 2 NaOH(aq) : Na2A(aq) + 2 H2O(𝓵)

If 36.04 mL of 0.509 M NaOH is required to titrate the acid to the equivalence point, what is the molar mass of the acid? 4.64

In this chapter, we used a simplified equation for the combustion of gasoline. But the products generated in the operation of real automobile engines include various oxides of nitrogen, and it is desirable to remove these compounds from the exhaust stream rather than release them into the atmosphere. Currently this is done by the “catalytic converter,” but other technologies are being explored. One possible method involves the use of isocyanic acid, HNCO, which reacts with NO2 to produce N2, CO2, and H2O, as shown in the equation below: ■

the reaction below, what mass of the mineral fluoroapatite (Ca5(PO4)3F) would be required each year? Ca5(PO4)3F + 5 H2SO4 : 3 H3PO4 + 5 CaSO4 + HF 4.66

■ The reaction shown below is used to destroy Freon-12 (CF2Cl2), preventing its release into the atmosphere. What mass of NaF will be formed if 250.0 kg of CF2Cl2 and 400.0 kg of Na2C2O4 are heated and allowed to react to completion?

CF2Cl2 + 2 Na2C2O4 : 2 NaF + 2 NaCl + C + 4 CO2 4.67 Blood alcohol levels are usually reported in mass percent: a level of 0.10 means that 0.10 g of alcohol (C2H5OH) was present in 100 g of blood. This is just one way of expressing the concentration of alcohol dissolved in the blood. Other concentration units could be used. Find the molarity of alcohol in the blood of a person with a blood alcohol level of 0.12. (The density of blood is 1.2 g/mL.) 4.68 One way of determining blood alcohol levels is by performing a titration on a sample of blood. In this process, the alcohol from the blood is oxidized by dichromate ions (Cr2O72−) according to the following net ionic equation: C2H5OH + 2 Cr2O72− + 16 H+ : 2 CO2 + 4 Cr3+ + 11 H2O A 10.00-g sample of blood was drawn from a patient, and 13.77 mL of 0.02538 M K2Cr2O7 was required to titrate the alcohol. What was the patient’s blood alcohol level? (See the previous problem for definition of blood alcohol level. K2Cr2O7 is a strong electrolyte, so it dissociates completely in solution.) 4.69 Ammonium sulfate ((NH4)2SO4) is a common fertilizer that can be produced by the reaction of ammonia (NH3) with sulfuric acid (H2SO4). Each year about 2 × 109 kg of ammonium sulfate is produced worldwide. How many kilograms of ammonia would be needed to generate this quantity of (NH4)2SO4? (Assume that the reaction goes to completion and no other products formed.) 4.70 The pictures below show a molecular scale view of a chemical reaction between the compounds AB2 and B2. (Green balls represent B atoms and orange balls are A atoms). The box on the left represents the reactants at the instant of mixing, and the box on the right shows what is left once the reaction has gone to completion.

8 HNCO + 6 NO2 : 7 N2 + 8 CO2 + 4 H2O What mass of HNCO would be required to react completely with the estimated 1.7 × 1010 kg of NO2 produced each year by U.S. automobiles? How might the fact that the reaction above produces CO2 affect its feasibility as a catalytic converter? 4.65

122

Phosphoric acid (H3PO4) is important in the production of both fertilizers and detergents. It is distributed commercially as a solution with a concentration of about 14.8 M. Approximately 2.1 × 109 gallons of this concentrated phosphoric acid solution is produced annually in the U.S. alone. Assuming that all of this H3PO4 is produced by ■


Was there a limiting reagent in this reaction? If so, what was it? Write a balanced chemical equation for this reaction. As usual, your equation should use the smallest possible whole number coefficients for all substances.

■


4.71 The pictures below show a molecular scale view of a chemical reaction between H2 and CO to produce methanol, CH3OH. The box on the left represents the reactants at the instant of mixing, and the box on the right shows what is left once the reaction has gone to completion.

4.77

■ Calcium

carbonate (limestone, CaCO3) dissolves in hydrochloric acid, producing water and carbon dioxide as shown in the following unbalanced net ionic equation. CaCO3(s) + H3O+(aq) : H2O(𝓵) + CO2(g) + Ca2+(aq) Suppose 5.0 g of CaCO3 is added to 700. mL of 0.10 M HCl. What is the maximum mass of CO2 that could be formed? What would the final concentration of Ca2+ ions be? (Assume that the final solution volume is still 700 mL.)

4.78

■ Iron metal can be refined from the mineral hematite (Fe2O3). One way of converting the mineral to iron is to react it with carbon monoxide, as shown below:

Fe2O3 + 3 CO : 2 Fe + 3 CO2

Was there a limiting reagent in this reaction? If so, what was it? Write a balanced chemical equation for this reaction. As usual, your equation should use the smallest possible whole number coefficients for all substances.

Because the hematite is obtained from various ores, it is usually not in a pure form. Suppose an iron manufacturer has 2.00 × 105 kg of ore available, and the ore is 93% Fe2O3 by mass. (There is no iron in the remaining 7% of the ore.) How many moles of Fe2O3 are present in this ore? How many kg of pure iron could be obtained from this sample of ore? Assume that the process has a 100% yield and that excess CO is available. 4.79

4.72 Aluminum chloride (AlCl3) is used as a catalyst in the production of polyisobutylene, which is used in automobile tires. Scrap aluminum metal reacts with chlorine gas (Cl2) to produce AlCl3. Suppose that 2.70 g of Al and 7.10 g of Cl2 are mixed. What is the maximum mass of AlCl3 that could be formed? 4.73 In the cold vulcanization of rubber, disulfur dichloride (S2Cl2) is used as a source of sulfur atoms, and those sulfur atoms form “bridges,” or cross-links, between polymer chains. S2Cl2 can be produced by reacting molten sulfur (S8(𝓵)) with chlorine (Cl2(g)). What is the maximum mass of S2Cl2 that can be produced by reacting 32.0 g of sulfur with 71.0 g of chlorine? 4.74

■ Styrene,

the building block of polystyrene, is a hydrocarbon. If 0.438 g of the compound is burned and produces 1.481 g CO2 and 0.303 g H2O, what is the empirical formula of the compound?

4.75

■ Cryolite

(Na3AlF6) is used in the commercial production of aluminum from ore. Cryolite itself is produced by the following reaction: 6 NaOH + Al2O3 + 12 HF : 2 Na3AlF6 + 9 H2O

A mixture containing 800.0 kg of NaOH, 300.0 kg of Al2O3, and 600.0 kg of HF is heated to 950°C and reacts to completion. What is the maximum mass of Na3AlF6 formed? 4.76 A quality control technician needs to determine the percentage of arsenic (As) in a particular pesticide. The pesticide is dissolved and all of the arsenic present is converted to arsenate ions (AsO43−). Then the amount of AsO43− is determined by titrating with a solution containing silver ions (Ag+). The silver reacts with the arsenate according to the following net ionic equation:

■


sulfate is the essential component of plaster and sheet rock. Waste calcium sulfate can be converted into quicklime, CaO, by reaction with carbon at high temperatures. The following two reactions represent a sequence of reactions that might take place: CaSO4(s) + 4 C(s) : CaS(𝓵) + 4 CO(g) CaS(𝓵) + 3 CaSO4(s) : 4 CaO(s) + 4 SO2(g) What mass of sulfur dioxide (in grams) could be obtained from 1.250 kg of calcium sulfate?

FOCUS ON PROBLEM SOLVING EXERCISES 4.80 The carbon dioxide exhaled by astronauts must be removed from the air. One way to do this is through the use of CO2 “scrubbers” based on metal hydroxides. CO2 reacts with the hydroxide to form the corresponding metal carbonate and water. Because these scrubbers need to be sent into space, weight is a significant concern in their design. If the goal is to minimize the mass of metal hydroxide needed to remove a given amount of CO2 from the air, which would you recommend using: sodium hydroxide, magnesium hydroxide, or aluminum hydroxide? Explain your answer using arguments based on stoichiometry. 4.81 You are designing a process for a wastewater treatment facility intended to remove aqueous HCl from the input water. Several bases are being considered for use, and their costs are listed in the table below. If cost is the primary concern, which base would you recommend using? Use stoichiometric arguments to justify your answer.

3 Ag+(aq) + AsO43−(aq) : Ag3AsO4(s) When a 1.22-g sample of pesticide was analyzed this way, it required 25.0 mL of 0.102 M Ag+ solution to precipitate all of the AsO43−. What was the mass percentage of arsenic in the pesticide?

■ Calcium

4.82

Material

CaCO3

Ca(OH)2

NH3

NaOH

Cost per metric ton

$270

$165

$255

$315

■ You are responsible for the design of a plant to produce silicon carbide, SiC, for use as an abrasive in grinding applications. The reaction you will use combines silicon dioxide with graphite (solid carbon) to form silicon carbide


123

and carbon monoxide. The silicon dioxide is to be obtained from sand, and the carbon is derived from coal. If the available sand is 95% SiO2 by weight and coal is 73% C by weight, what mass of coal is needed for each metric ton of sand used? 4.83 Chromium metal is found in ores as an oxide, Cr2O3. This ore can be converted to metallic Cr by the Goldschmidt process, Cr2O3(s) + 2 Al(s) : 2 Cr(𝓵) + Al2O3(s). Suppose that you have 4.5 tons of chromium(III) oxide. Describe how you would determine how much aluminum metal you would need to carry out this process. List any information you would have to look up.

dioxide, describe how to determine what mass of CO2 is evolved during the process. 4.91 Existing stockpiles of the refrigerant Freon-12, CF2Cl2, must be destroyed under the terms of the Montreal Protocol because of their potential for harming the ozone layer. One method for doing this involves reaction with sodium oxalate: CF2Cl2 + 2 Na2C2O4 : 2 NaF + 2 NaCl + C + 4 CO2 If you had 150 tons of Freon-12, describe how you would know how much sodium oxalate you would need to make that conversion.

4.84 Hydrochloric acid reacts with many metals to form hydrogen gas, H2. If you needed to produce hydrogen as a source of fuel in some design, how would you determine which metal, Mg, Al, or Zn, would produce the greatest amount of hydrogen per gram of metal? Describe the steps you would need to take in the decision making process. 4.85 When heated in air, some element, X, reacts with oxygen to form an oxide with the chemical formula X2O3. If a 0.5386-g sample of this unknown element yields an oxide with a mass of 0.7111 g, describe how you would determine the atomic mass of the element.

Cumulative Problems

4.86 Chlorofluorocarbons were used for many years as refrigerants. One such molecule, C2Cl2F2, is synthesized from the reaction of CCl4 with HF. Suppose that you have 150 tons of carbon tetrachloride as a starting point for this reaction, show how you would determine the minimum amount of HF needed to carry out the reaction.

4.92 Elemental analysis is sometimes carried out by combustion of the sample. For a hydrocarbon, the only products formed are CO2 and H2O. If a 1.36-g sample of an unknown hydrocarbon is burned and 2.21 g of H2O is produced along with 4.07 g of CO2, what is the empirical formula of the hydrocarbon?

4.87 Bismuth selenide, Bi 2Se 3, is used in semiconductor research. It can be prepared directly from its elements by the reaction, 2 Bi + 3 Se : Bi2Se3. Suppose that you need 5 kg of bismuth selenide for a prototype system you are working on. You find a large bottle of bismuth metal but only 560 g of selenium. Describe how you can calculate if you have enough selenium on hand to carry out the reaction or if you need to order more.

4.93 Benzopyrene is a hydrocarbon that is known to cause cancer. Combustion analysis (see Problem 4.92) of a 2.44-g sample of this compound shows that combustion produces 8.54 g of CO2 and 1.045 g of H2O. What is the empirical formula of benzopyrene?

4.88 When 2.750 g of the oxide of lead Pb3O4 is heated to a high temperature, it decomposes to produce 0.0640 g of oxygen gas and 2.686 g of some new lead oxide compound. How can you use this data to determine the formula of the new compound? 4.89 The best way to generate small amounts of oxygen gas in a laboratory setting is by decomposing potassium permanganate. 2 KMnO4(s) : K2MnO4 + MnO2 + O2 Suppose that you are testing a new fuel for automobiles and you want to generate specific amounts of oxygen with this reaction. For one test, you need exactly 500 g of oxygen. Describe how you would determine the mass of potassium permanganate you need to obtain this amount of oxygen. 4.90 An ore sample with a mass of 670 kg contains 27.7% magnesium carbonate, MgCO3. If all of the magnesium carbonate in this ore sample is decomposed to form carbon

124


Freon-12, CF2Cl2

4.94 Write the balanced chemical equation for the combustion of ethane, C2H6, and answer these questions. (a) How many molecules of oxygen would combine with 14 molecules of ethane in this reaction? (b) If 13.0 mol of oxygen is consumed in a reaction, how many moles of water are produced? (c) How many grams of ethane are burned if 4.20 × 1022 molecules of CO2 are produced? 4.95 Aluminum metal reacts with sulfuric acid to form hydrogen gas and aluminum sulfate. (a) Write a balanced chemical equation for this reaction. (b) Suppose that a 0.792-g sample of aluminum that contains impurities is reacted with excess sulfuric acid and 0.0813 g of H2 is collected. Assuming that none of the impurities reacts with sulfuric acid to produce hydrogen, what is the percentage of aluminum in the sample? 4.96 A metallurgical firm wishes to dispose of 1300 gallons of waste sulfuric acid whose molarity is 1.37 M. Before disposal, it will be reacted with calcium hydroxide (slaked lime), which costs $0.23 per pound. (a) Write the balanced chemical equation for this process. (b) Determine the cost that the firm will incur from this use of slaked lime.

■


5

Gases

OUTLINE 5.1

INSIGHT INTO Air Pollution

5.2

Pressure

5.3

History and Application of the Gas Law

5.4

Partial Pressure

5.5

Stoichiometry of Reactions Involving Gases

5.6

Kinetic–Molecular Theory and Ideal Versus Real Gases

5.7

INSIGHT INTO Gas Sensors

Urban smog results from a complex sequence of chemical reactions in the atmosphere. Knowledge of the properties of gases is essential for a meaningful understanding of smog and other issues involving air pollution. NASA

I

n many ways, the essence of chemistry is considering things on the atomic and molecular level. This idea applies when we study chemical reactions and also when we examine the properties of matter. So as we begin to look at matter and materials from a chemist’s perspective, we will always want to focus on the way the properties of matter are related to the properties and behavior of its constituent molecules. Gases provide an excellent starting point for this because the observable macroscopic properties of gases are very direct results of the behavior of individual molecules within the gas. Most of the earliest investigations of gas properties were carried out using air. The reasons for this are clear: air is both readily available and obviously important to our lives. Although the sophistication of the measurements we make has increased tremendously, the availability and importance of air have ensured that it remains a very carefully studied gas. Nowadays, much of our interest in air is actually driven by concerns over the levels and effects of various pollutants. The U.S. Environmental Protection Agency (EPA) defines pollutants as “unwanted chemicals or other materials found in the air.” Although this definition does include a variety of small solid particles such as dust or soot, most air pollutants are gases. Thus an exploration of some of the issues surrounding the monitoring and control of air pollution will form the backdrop for our study of gases and their properties.


125

Chapter Objectives After mastering this chapter, you should be able to ❚ describe the physical properties of gases. ❚ identify several gaseous compounds or classes of compounds that are important in urban air pollution. ❚ use the ideal gas law for calculating changes in the conditions of a gas. ❚ use the concept of partial pressure to work with mixtures of gases. ❚ perform stoichiometric calculations for reactions involving gases as reactants or products. ❚ state the postulates of the kinetic theory of gases. ❚ describe qualitatively how the postulates of the kinetic theory account for the observed behavior of gases. ❚ describe the Maxwell-Boltzmann distribution of speeds and the effects of temperature and molar mass on molecular speed. ❚ identify conditions under which gases might not behave ideally. ❚ use the van der Waals equation to perform calculations for gases under nonideal conditions. ❚ describe the principles of operation for some pressure-measuring devices.

INSIGHT INTO

5.1

Air Pollution

Air pollution is one of the unintended but unavoidable side effects of modern society. The issues surrounding the monitoring, prevention, and remediation of air pollution involve an exceedingly complex interplay between science, business, and public policy. Our discussion here will be fairly limited, but we will try to explore some of the ways in which chemical science can help us to understand the complexity of the situation. We should begin by looking at the chemical composition of clean air. As you probably know, air is a mixture rather than a pure substance. Nitrogen and oxygen account for the majority of the molecules present. Other important components include carbon dioxide and water vapor. The amount of water in the air varies from place to place and day to day, as described by the humidity in a weather report. So it will be convenient to use dry air as a reference point in our discussion. Table 5.1 lists the composition of a one cubic meter sample of dry air at 25°C and normal atmospheric Table

❚

5.1

The composition of a one cubic meter sample of dry air at 25°C and normal atmospheric pressure

Gas N2

31.929

O2

8.567

Ar

0.382

CO2

0.013

Other trace gases

0.002

Total

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Chapter 5 Gases

# Moles Present

40.893

pressure. The four gases listed account for at least 99.99% of the molecules present; N2 and O2 alone make up roughly 99%. In contrast to the short list in Table 5.1, if we were to collect a sample of air in a modern urban area and analyze it carefully, we might detect dozens of compounds at trace levels. Some of these additional trace components can be linked to natural sources, but most are directly attributable to human activities like transportation and manufacturing. Of the large number of chemical compounds found in urban air, the EPA has identified six principal criteria pollutants: CO, NO2, O3, SO2, Pb, and particulate matter (PM). These six substances are commonly found throughout the country, causing a variety of negative effects on health, the environment, and property. The term “criteria pollutants” comes from the fact that the EPA has established a set of science-based standards, or criteria, for the acceptable levels of these pollutants. Primary standards are intended to protect our health, and secondary standards are intended to protect our environment and property. In many cases, the allowable levels of these pollutants correspond to less than one part per million, which means one molecule per every million molecules of air. Levels of these pollutants are monitored periodically throughout the country. If the detected concentrations exceed the primary standard in a given region, that area is designated as a nonattainment area for that particular pollutant and additional regulatory action may be taken. As an example to illustrate the complexity of pollution issues, let’s consider one of the criteria pollutants, NO2. We might begin by asking where this pollutant comes from. Nitrogen dioxide is emitted as part of the exhaust gas from our automobiles. At the high temperatures at which engines operate, the nitrogen and oxygen gases present in air react with each other to form a variety of nitrogen oxides. Such reactions do not occur at lower temperatures but are practically unavoidable in an engine. When engines run hotter, from standing in a traffic jam, for example, they tend to produce even more nitrogen oxides. Once it is present in the air, NO2 poses a health risk because it attacks the membranes in our lungs. The brown color of smog is also largely due to NO2. Nitrogen dioxide molecules absorb some portion of the visible light of the Sun, leaving the colors that produce brown to pass through. The detrimental effects of NO2 in the air extend well beyond this rather unappealing color. Light energy from the Sun can also initiate chemical changes through what is known as a photochemical reaction. For nitrogen oxides, these photochemical processes lead to products that can react further to trigger the formation of ground-level ozone. But nitrogen oxides cannot generate ozone on their own. The key additional components that are needed are molecules collectively referred to as volatile organic chemicals, or VOCs. As we learned in Section 2.6, organic chemicals contain predominantly carbon and hydrogen. Volatile compounds tend to evaporate easily. VOCs combine these two traits, meaning they are carbon-based molecules that evaporate readily. Many different VOCs are present in urban air. The primary source of these chemicals is fuels. In Section 4.1, we learned that the gasoline we burn in our automobiles and lawn mowers is a complex mixture of organic molecules. All combustion engines release some VOCs when they are operating, and older vehicles tend to release more. Our noses detect this phenomenon every time we are near a vehicle that produces large amounts of VOCs. Chemical reaction between photochemically activated nitrogen oxides and VOCs can give rise to a mixture of gases that is collectively referred to as smog. Many of the compounds in smog are irritants to human lungs. Ozone (O3) is the most significant problem. When present at ground level, ozone poses a range of health hazards, even for healthy young adults. People with respiratory problems, such as asthma, are particularly susceptible to problems associated with ozone. Figure 5.1 shows a graph of concentrations of various chemicals associated with photochemical smog as a function of time of day in a large city. Notice that the nitrogen oxides and VOCs are both present prior to the formation of ozone. If either of these molecules is missing, ozone will not form. As we will see, this observation provides a way to combat the formation of

Particulate matter dispersed in the air includes soot, dust, and other small solid particles.

The various oxides of nitrogen are often referred to collectively as “NOX.”

We will look at the desirable effects of ozone in the upper atmosphere in Chapter 11.

5.1

Air Pollution

127

0.14 NO

Concentration (ppm)

0.12 0.10

NO2

Lung irritants (primarily ozone)

0.08 0.06 0.04 VOC 0.02 0 Midnight

3 am

6 am

9 am Noon 3 pm Time of day (hour)

6 pm

9 pm

Midnight

Figure 5.1 ❚ The graph shows how the levels of various air pollutants in an urban area vary during a day. The levels of nitrogen oxides and VOCs rise in the morning hours, mainly due to emissions from vehicles. The presence of both NOx and VOCs leads to increased production of lung irritants such as ozone, whose levels go through a maximum during the afternoon hours.

ozone. Many cities have adopted programs that target one or more of the key constituents in photochemical smog, hoping to avoid the formation of its most hazardous components. Even in this brief overview, we see that a single pollutant—NO2 in this case—can have a complex impact on the air around us. We also see that the effects of various pollutants can be coupled in unexpected ways; the simultaneous presence of NO2 and VOCs leads to the generation of O3, another of the criteria pollutants. Clearly, the chemistry of air pollution is quite complex, so we will not be able to delve into many aspects of it in great detail. But it should be clear that to improve our understanding of these important issues, we must first explore the properties of gases, gas mixtures, and chemical reactions involving gases.

Properties of Gases Several observable properties are common to all gases and distinguish them from solids and liquids. • • • • •

Gases expand to fill the volume of any container. Gases have much lower densities than solids or liquids. Gases have highly variable densities, depending on the conditions. Gases mix with one another readily and thoroughly. Gases change volume dramatically with changing temperature.

Some of these properties help explain why air pollution is difficult to combat. The fact that gases expand to fill the available volume, for example, means that exhaust gases released from a car or an industrial plant will readily spread throughout the surrounding area. Similarly, the fact that gases mix readily with one another means that once they are present, traces of gaseous pollutants will usually be difficult to separate from the air. As we work toward developing an accurate model for gas behavior, we will want to be sure that it can account for all of these observations. We’ll want our model to be based on a molecular level understanding of what a gas is like, and we’ll also want it to enable us to make accurate numerical predictions about how a gas will respond to different conditions.

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As it turns out, you probably already know an important quantitative part of the model we use most often to describe gases. If there is just one equation that you remember from your high school chemistry class, chances are it is the ideal gas law: PV = nRT

(5.1)

P is the gas pressure, V is the volume, n is the number of moles of gas present, and T is the temperature. R is a proportionality constant, called the universal gas constant because its value (0.08206 L atm mol−1 K−1) is the same for any gas. This familiar equation will form the basis of much of our discussion in this chapter, so it will be important that we have a thorough understanding of the terms that appear in it. Let’s start by taking a close look at the meaning and physical origin of gas pressure.

5.2

R can also be expressed in other units, as discussed in Section 5.3.

Pressure

From inflation guidelines for car or bicycle tires to weather reports, we regularly encounter the word pressure, so you undoubtedly have some notion of what it means. But to investigate the behavior of gases scientifically, we need to make our intuitive understanding of pressure more formal. Pressure is defined as force exerted per unit area: F P=— A

(5.2)

The obvious question arising from this definition is, “What force causes the pressure of a gas?” For atmospheric pressure, the force can be attributed to the weight of the molecules of air that are attracted to the surface of the planet by gravity. The existence of atmospheric pressure may not be apparent but can readily be demonstrated by removing the air from a container. Even an apparently sturdy container collapses under the pressure of the atmosphere when the air inside is no longer supporting the container walls. Because the total weight of the molecules increases as the number of molecules increases, the atmospheric pressure changes as a function of altitude. At sea level, 20,320 feet more of atmosphere are above us than at the summit of Mount McKinley, and the atmospheric pressure at sea level is much higher than at the mountaintop. The variation of pressure with altitude is shown in Figure 5.2. But what about gas trapped in a container? We can think about this by considering the air in a balloon. It exerts pressure uniformly on all the walls of the balloon without regard to direction. Clearly, this can’t be due to gravity. How does a gas exert force on the walls of its container? To answer this, we need to zoom in to the molecular level, as

16

Figure 5.2 ❚ Atmospheric pressure varies as a function of altitude. Pressure is given here in units of pounds per square inch. Other units of pressure are discussed later in this section.

14

Pressure (psi)

12 10 8 6 4 2 0

20,000

40,000

60,000 Altitude (feet)

80,000

100,000

120,000

5.2

Pressure

129

Figure 5.3 ❚ The inset in this


figure shows a microscopic scale “close-up” of the gas molecules in a balloon. As molecules move through the volume of the balloon, occasionally they collide with the rubber walls. In every such collision, the gas molecules exert a small force on the balloon wall, pushing it outward. When summed over the large number of gas molecules present, these small forces give rise to the observed pressure of the gas and are responsible for the shape of the balloon.

in Figure 5.3, where we will see individual gas molecules colliding with the walls and bouncing off. Each molecular collision imparts a small force to the container. And when these tiny forces are summed across the vast number of gas molecules present, they can produce the macroscopic force that leads to gas pressure. Later we will see how several observable gas properties can readily be explained by considering the behavior of gas molecules on the microscopic scale.

Measuring Pressure

Scientists use the word ‘vacuum’ to refer to any region of reduced pressure, not just the idealized case of zero pressure.

Because pressure is one of the important measurable properties of a gas, several different ways have been devised to measure it, beginning with the barometer. The essential feature of a standard barometer is a long tube that is closed at one end and is filled with mercury. (Mercury is used because it has a high density and does not react chemically with common gases.) When the mercury-filled tube is inverted in a pool of mercury, some of the liquid, but not all, runs out of the tube. The closed end of the tube is then under vacuum, which means that the pressure in this region is greatly reduced. If this end of the tube is opened, all of the mercury pours out. What prevents the movement of mercury when the tube is closed? We can understand how the pressure of air is responsible by referring to Figure 5.4. The pool of mercury is exposed to the atmosphere and thereby experiences the pressure of air weighing down on it. The column of mercury in the tube has a vacuum over it and so has no pressure weighing down on it. Now consider a point at the mouth of the tube in the pool of mercury. This point experiences pressure from two places, the atmospheric pressure on the pool and the weight of the mercury from the column. When these two pressures are equal, the height of mercury suspended in the tube remains constant and provides a way to determine the atmospheric pressure. When the atmospheric pressure decreases (from weather variations) the decrease will cause the column to shorten to maintain the equality of pressure at that point. Near sea level, the column of mercury is typically between 730 mm and 760 mm in height. We’ll look at some more modern ways of measuring pressure in the closing section of this chapter.

Units of Pressure The SI unit for pressure is a derived unit that reflects the definition of pressure. The SI units for force and area are the newton and the square meter, respectively. So pressure is measured in N m−2. This unit has also been named the pascal, Pa. One newton is roughly the force exerted by gravity on a one-quarter-pound object. When spread

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Chapter 5 Gases

Figure 5.4 ❚ A schematic diagram 800 700 600 500 400

Force of mercury

300 Force of 200 atmosphere

of the forces at work in a mercury barometer. Atmospheric pressure exerts a force pushing down on the surface of the mercury and tends to push the column of mercury in the barometer higher up the tube. The force due to the weight of the mercury in the column tends to push the column of mercury lower in the tube. When the two forces are equal, the column height remains constant.

Mercury

100 0

out over a square meter, this force does not produce very much pressure, so the Pa is a relatively small quantity. Typical atmospheric pressures are on the order of 105 Pa. Thus a more practical unit is kPa, which is the unit used in weather forecasts in Canada and other “metric countries.” The construction of the barometer has influenced the units commonly used for reporting pressures. Most weather forecasts in the United States report pressure in inches (inches of mercury). Most barometers in school laboratories provide measurements in millimeters of mercury (mm Hg). Both of these units appear to express pressure as a length, which can be confusing because we defined pressure as force per unit of area. The explanation of this contradiction comes directly from the barometer. The height of the mercury column in a barometer is directly proportional to the applied pressure; so the higher the pressure exerted by the atmosphere, the higher the column of mercury. As a result, it became common to report pressure directly in terms of the measured height. In spite of this confusion, the mm Hg unit is still widely used and has been given an additional name, the torr. We can consider these two units equivalent: 1 torr = 1 mm Hg Another unit of pressure arises from the importance of the atmosphere. The average pressure at sea level is close to 760 torr. The common observation of an apparent average atmospheric pressure led to the definition of a unit called an atmosphere (atm), which was set equal to 760 torr. As science has become more precise in measuring it, this equivalence still remains:

The torr is named after Evangelista Torricelli (1608–1647), who was the first person to suggest building a mercury barometer.

1 atm = 760 torr (exactly) But the definition of the unit atm has since been changed to reflect SI units. It is now defined in terms of the pascal: 1 atm = 101,325 Pa (exactly) Combining the last two relationships gives us a formal definition of the torr as well: 760 torr = 101,325 Pa (exactly) The three units that we will use most often in this book are the atm, torr, and kPa. In some engineering fields, you will also be likely to encounter other units, including

5.2

Pressure

131

pounds per square inch. None of these should cause you particular difficulty as long as you are careful that any calculations you do properly account for the units used. Now, we might ask how gas pressure plays a role in air pollution. Gases naturally flow from higher pressure to lower pressure, and pollutants are often produced in a high-pressure environment, such as in the cylinder of an auto engine or a reaction vessel in an industrial processing plant. When gases are released from these high-pressure environments, they disperse rapidly into the surrounding air. The pressure difference helps speed the spread of the pollutants.

5.3

History and Application of the Gas Law

Unlike solids and liquids, gases change significantly when the conditions in which they are found are altered. Changing the temperature, pressure, or volume of a gas’s container results in a dramatic response from the gas. All gases respond to these physical changes in the same way despite their chemical differences, and the gas law equation allows us to calculate the impact of a particular change. Historically, the study of gases preceded most of the important discoveries of modern chemistry. The relationships between physical properties of gases summarized by the ideal gas law equation were initially based on empirical observations and often still carry the name of the investigator who first characterized the relationship. Jacques Charles first investigated the relationship between temperature and volume in the early 19th century. He noted that when he plotted volume versus temperature for different amounts of gas, the resulting lines converged to the same temperature at zero volume (see Figure 5.5). This fact was used to establish the Kelvin temperature scale. Once the temperature is expressed on the Kelvin scale, the volume and temperature are directly proportional, an observation now known as Charles’s law. Other early scientists examined the connections between pairs of gas variables and arrived at gas laws to express their observations. Robert Boyle studied the relationship between gas pressure and volume and noted that for any gas sample, when pressure increases, volume decreases, or when volume increases, pressure decreases. Thus Boyle’s law states that at constant temperature, pressure and volume are inversely proportional. Similarly, Avogadro’s law states that the volume of a gas is proportional to the number of molecules (or moles) of gas present. These laws are frequently used to calculate the effect of some particular change in gas conditions. Fortunately, we will not need to memorize specific equations for all of these individual laws because they are all incorporated into the ideal gas law. As long as we know the ideal gas equation and know how to work with it, we can easily solve problems that might also be approached using any of the individual gas laws. To do this, just start by identifying all of the parameters from the gas law—P, V, T, or n—that

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Chapter 5 Gases

50

40

Volume

Figure 5.5 ❚ The graph shows volume as a function of temperature for two different samples of gas. The sample represented by the upper red line contains more gas molecules than that represented by the lower blue line. Note that for both samples, the relationship between volume and temperature is clearly linear. Also note that when both lines are extrapolated back toward lower temperatures, they have the same x intercept at −273.15°C. This extrapolation defines the so-called “absolute zero” temperature, or 0 K.

30

20

Absolute zero –273°C

10

0 –300

–200

–100

0 Temperature (˚C)

100

200

300

do not change. (R is a universal constant, so it never changes!) Then rearrange the gas law equation to collect all of those terms onto the same side, leaving those that do change on the other side. If pressure and volume change while the number of moles and the temperature remain constant, for example, we’d get PV = nRT = constant Because the product nRT remains constant, we know that PV must also be constant. This tells us that the product of P times V must be the same before and after the change, so we can write P1V1 = P2V2

(5.3)

Here the subscripts ‘1’ and ‘2’ indicate the conditions before and after the change takes place. Now if we know any three of those four quantities, we can solve for the missing one, as seen in the following example.


5 .1

In addition to ongoing sources, such as car engines, pollutants can also be introduced into the air through isolated incidents like the rupture of a gas storage tank. Most gases are stored and transported at high pressures. A common laboratory cylinder of methane, for example, has a volume of 49.0 L and is filled to a pressure of 154 atm. Suppose that all of the CH4 from this cylinder is released and expands until its pressure falls to 1.00 atm. What volume would the CH4 occupy?

Strategy The problem involves a change in the conditions, here pressure and volume, of a sample of gas. We will want to use the ideal gas law because it is our model that relates the various properties of a gas to one another. So we begin by assuming that the gas will obey the ideal gas law under both the initial and final conditions. We will further assume that temperature is constant because we are not given any data to suggest otherwise. Because we are working with the same sample of gas throughout, we also know that the number of moles (n) is constant. Collect the constant terms together and work toward a solution.

In Section 5.6, we will look at gases that do not behave ideally.

Solution In this case, n and T are constant, and R is always constant. So, PV = nRT = constant Therefore, P1V1 = P2V2 Solving this for the final volume (V2) gives P1V1 V2 = — P2 We know all three terms on the right: P1 = 154 atm, V1 = 49.0 L, and P2 = 1.00 atm. Insert those above and solve: (154 atm)(49.0 L) V2 = — = 7550 L (1.00 atm)

Analyze Your Answer Our result is in liters, which is an appropriate unit for volume. The actual value is much larger than the initial volume, so we might ask if this seems sensible. Looking at the initial and final conditions, we have lowered the pressure by a large factor, and this should lead to a much larger volume. So our answer seems plausible.

Discussion Because of the high initial pressure, the assumption that the gas behaves ideally may not be very accurate. We’ll examine that assumption later in Example 5.8. 5.3 History and Application of the Gas Law

133

Check Your Understanding A small sample of gas is generated in a laboratory. In a 125-mL vessel, it exerts 115 torr of pressure. Then a valve is opened and the gas expands until its volume is 175 mL. If the temperature is constant throughout this process, what is the new pressure of the gas?

In other situations, we might find that V and T vary, whereas n and P are constant. That would lead us to write the following relationship V T

nR P

— = — = constant

This in turn leads to Equation 5.4: V1 T1

V2 T2

—=—

E XAM PL E PRO BLEM

(5.4)

5 .2

A balloon is filled with helium, and its volume is 2.2 L at 298 K. The balloon is then dunked into a thermos bottle containing liquid nitrogen. When the helium in the balloon has cooled to the temperature of the liquid nitrogen (77 K), what will the volume of the balloon be?

Strategy We should start by recognizing this as a change in conditions for a gas and realize that we will want to use the ideal gas law. So we assume that the gas will behave ideally under both the initial and final conditions. We will further assume that the pressure is constant since we are not given any data to suggest otherwise. And since the balloon is tied off throughout, we can also assume that n is constant. Collect the constant terms together and work toward a solution.

Solution As in the discussion above, we have V T

nR P

— = — = constant

So V1 T1

V2 T2

—=—

Solve this for the final volume (V2): V1T2 V2 = — T1 Because the three terms on the right are known, we can just insert values and solve: (2.2 L)(77 K) V2 = — = 0.57 L (298 K)

Analyze Your Answer We have decreased the temperature from 298 K to 77 K, which means that it is smaller by a factor of nearly 4. This should lead to a volume that is also smaller by a factor of nearly 4. So our answer seems likely to be correct.

Check Your Understanding The balloon in the example above will burst if its volume exceeds 2.3 L. At what temperature would you expect the balloon to burst?

For any change in the conditions of a gas sample, we should always be able to employ this same idea by separating those parameters that are constant from those that 134

Chapter 5 Gases

are changing. Because you are very unlikely to forget the ideal gas law, this approach is almost certainly more reliable than attempting to remember a list of equations covering all the possible combinations of variables.

Units and the Ideal Gas Law Suppose that we had expressed the temperature in °C rather than in kelvins in the above example? It should take you only a little work to convince yourself that the equation we used will not hold true. Changing the temperature of a gas from 10 to 20°C, for example, will not double its volume. The ideal gas law itself is only valid when we use an absolute temperature scale where T = 0 corresponds to absolute zero temperature. To see why this must be so, notice that putting T = 0 into the gas law implies that either P or V or both must also go to zero, meaning that there must be no gas present. This simply can’t be correct for T = 0°C or T = 0°F; that’s obvious because the winter temperature routinely reaches these levels in many areas, and yet the air in the atmosphere obviously does not vanish. Absolute zero (0 K), on the other hand, is a very special temperature, and it is reasonable that there might be peculiar consequences if we actually cooled a gas to this level. In contrast to the temperature, which must be expressed on an absolute scale, we can use any convenient units for pressure and volume, so long as we then reconcile the units on those quantities with those for the gas constant, R. Why this difference? The common temperature scales are fundamentally different from other kinds of units because converting among them involves not only multiplying by a conversion factor but also adding or subtracting. This, in turn, results directly from the fact that different temperature scales (K, °C, or °F ) have different zero points. In all the other quantities that we work with, the meaning of “zero” is the same no matter what units we use. If you say the mass of a sample is zero, for example, it doesn’t really matter much what units you measured in—there’s still nothing there. But for temperature, the differences among 0°C, 0°F, and 0 K are obvious. In some engineering fields, you may encounter the Rankine temperature scale (°R). This is also an absolute scale, meaning that 0°R = 0 K. But one degree on the Rankine scale is the same size as 1°F. Because they are both absolute scales, Rankine and Kelvin temperatures can be interconverted by using a multiplicative unit factor: 1°R = 1.8 K. If you look at a table like the one in Appendix B of this book, you will find values for R in several combinations of units.

Although absolute zero is a fairly abstract concept, temperatures well below 1 K can be achieved routinely in the laboratory.

No degree sign is used when writing Kelvin temperatures.

R = 0.08206 L atm mol−1 K−1 = 8.314 J mol−1 K−1 = 62.37 L torr mol−1 K−1 Students often ask how to decide which of these values is “the right R” to use in a given situation. The answer is that they are all really the same quantity, just expressed in different units. (If you are curious, you can do unit conversions to verify this for yourself.) If pressure is given in atm and volume in L, then it will be convenient to use R as 0.08206 L atm mol−1 K−1, for example. But if pressures are given in torr, it might be easier to insert 62.37 L torr mol−1 K−1 for R. Or, if you always prefer to use the same value of R, you could simply convert pressures from any other units into atm before doing any calculations. It doesn’t matter which set of units we choose, as long as we are careful to reconcile the units in all of our calculations.


5.3

A sample of CO2 gas has a volume of 575 cm3 at 752 torr and 72°F. What is the mass of carbon dioxide in this sample?

Strategy We are trying to find the mass of a sample of gas. If we knew how many moles of CO2 were in the sample, we could easily use the molar mass to find the mass 5.3 History and Application of the Gas Law

135

of the sample. And since the problem involves a gas, we can attempt to use the ideal gas law to find the number of moles. We are given values for V, P, and T, so we will be able to solve for n. Some of the units given are unusual, so we will need to be careful to handle them properly.

Solution We can begin by converting some of the given information into units more commonly used for gases. T = (72 − 32) × 5/9 = 22°C = 295 K V = 575 cm3 = 575 mL = 0.575 L For P, we will use the value in torr and choose the corresponding value for R. We could also convert from torr to atm, of course. (752 torr)(0.575 L) PV n = — = —— = 0.0235 mol RT (62.37 L torr mol−1 K−1)(295 K) The gas in our sample is CO2, with a molar mass of 44.0 g/mol, so 44.0 g 0.0235 mol × — = 1.03 g CO2 mol

Analyze Your Answer The temperature and pressure are near typical room conditions, and the volume is a little more than a half liter. So this should correspond to a small balloon filled with CO2. One gram is a pretty small mass but seems reasonable for the gas in a small balloon.

Check Your Understanding Find the volume in ft3 of 1.0 mole of an ideal gas at −25°C and 710 torr.

5.4

Partial Pressure

By using the ideal gas law, we can determine many properties of an individual gas. In many instances, however, we encounter interesting situations with more than one gas present. Our consideration of the levels of pollutants in air is an obvious example. Even if we ignore any pollutants present, clean air is already a mixture of gases. Aside from the possibility of chemical reactions, how do the observed properties of a gas mixture differ from those of pure gases? The answer to this question is actually suggested by the gas law. In particular, note that none of the terms in the gas law equation depends on the identity of the gas. Consider a sample of air and for simplicity assume that it consists entirely of N2 and O2. Assuming that the air behaves ideally, its pressure can be expressed using the gas law: nRT P=— V Here n is the total number of moles of air present. But this total number of moles will simply be the sum of the moles of each component: ntotal = nN + nO 2

2

We can insert this into the expression above, giving (nN + nO )RT nN RT nO RT nRT P=—=—=—+— V V V V 2

2

2

2

Each of those last two nRT/V terms has a special significance; they are the pressures that we would see if only one of the components of the gas mixture were

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Chapter 5 Gases

present. If the nitrogen were the only gas present, nO would be zero, and we would be left with 2

nN RT P = PN = — V 2

2

If only oxygen were present, we could write a similar result for PO . For a mixture, the observed pressure is the sum of the hypothetical partial pressures of the individual gases. This concept, first established by John Dalton, is sometimes referred to as Dalton’s law of partial pressures: The pressure of a mixture of gases is the sum of the partial pressures of the component gases. In simple terms, the partial pressure is what the pressure would be if only that component were present. As long as the component gases behave ideally and do not actually react with one another, this relationship will hold for any gas mixture. For an arbitrary gas mixture, we can write the partial pressure of each component as 2

niRT Pi = — V

John Dalton is also known for his model proposing the existence and nature of atoms.

(5.5)

where the index, i, indicates which gas we are talking about (for a mixture of four gases, i would vary from 1 to 4 yielding, P1, P2, P3, and P4) and ni is the number of moles of component i. The total pressure of the mixture, P, is the sum of these partial pressures: RT P = ∑ Pi = ∑ ni — i i V

(5.6)

The total number of moles is the sum of the numbers of moles of the individual components: RT P = ntotal — V

The upper case Greek letter sigma (Σ) is used to indicate the sum of a number of similar terms.

(5.7)

These expressions suggest another way to understand partial pressures. We might choose to express the concentration of each component gas in terms of its mole fraction, defined as ni Xi = — n

(5.8)

total

For a pure gas, ni = ntot and the mole fraction is one. Any mixture will have ni < ntot, so the value of the mole fraction has a range from zero to one. If we take the ratio of a partial pressure to total pressure, using Equations 5.5, 5.7, and 5.8, Pi

ni (RT/V )

ni

—=—=— ntotal = Xi n (RT/V ) P

(5.9)

total

So the partial pressure is given by the mole fraction times the total pressure. Pi = Xi P

(5.10)

This set of equations provides the numerical means for us to use partial pressures, as shown in Example Problems 5.4 and 5.5.


5.4

Not all pollution is due to human activity. Natural sources, including volcanoes, also contribute to air pollution. A scientist tries to generate a mixture of gases similar to those found in a volcano by introducing 15.0 g of water vapor, 3.5 g of SO2 and 1.0 g of CO2 into a 40.0-L vessel held at 120.0°C. Calculate the partial pressure of each gas and the total pressure.

5.4

Partial Pressure

137

Strategy As usual, we will begin by assuming that all the gases behave ideally. Then we can treat each gas separately to determine its partial pressure. Finally, sum these partial pressures to calculate the total pressure.

Solution For all of the gases, T = 120°C = 393 K, and V = 40.0 L. 1 mol H2O H2O: 15.0 g H2O × — = 0.833 mol 18.0 g H2O (0.833 mol)(0.08206 L atm mol−1 K−1)(393 K) PH O = ——— = 0.672 atm 40.0 L 2

1 mol SO2 SO2: 3.5 g SO2 × — = 0.055 mol 64.1 g SO2 (0.055 mol)(0.08206 L atm mol−1 K−1)(393 K) PSO2 = ——— = 0.044 atm 40.0 L 1 mol CO2 CO2: 1.0 g CO2 × — = 0.023 mol 44.0 g CO2 (0.023 mol)(0.08206 L atm mol−1 K−1)(393 K) PCO2 = ——— = 0.018 atm 40.0 L Total pressure is the sum of these: Ptot = PH2O + PSO2 + PCO2 = 0.672 + 0.044 + 0.018 = 0.734 atm Note that you could also have used the total number of moles to find Ptot and then used mole fractions to find the partial pressures.

Analyze Your Answer It is possible to gain a feeling for quantities such as pressure and number of moles if you work a lot of these problems. In this case, we have almost a mole of gas in a volume of 40 L, and the temperature is somewhat high. Our result of 0.7 atm is not unreasonable.

Check Your Understanding A mixture containing 13.5 g of oxygen and 60.4 g of N2 exerts a pressure of 2.13 atm at 25°C. What is the volume of the container, and what are the partial pressures of each gas?

E XAM PL E PRO BLEM

5 .5

An experiment is designed to determine the effect of sulfur dioxide, one of the EPA criteria pollutants, on plants. Among the variations used is a mixture that has the mole fractions given in the following table. Gas Mole fraction

N2

O2

H2O

SO2

0.751

0.149

0.080

0.020

If the desired total pressure is 750. torr, what should the partial pressures be? If the gas is to be in a 15.0-L vessel held at 30°C, how many moles of each substance are needed?

Strategy We know the mole fractions and the desired total pressure. So we can calculate partial pressures using the relationship defined in Equation 5.10. From the total pressure and the volume, we can calculate the total number of moles and thereby the number of moles of each gas.

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Solution Pi = X i P total P N = (0.751)(750. torr) = 563 torr 2

P O = (0.149)(750. torr) = 112 torr 2

P H O = (0.080)(750. torr) = 60 torr 2

P SO = (0.020)(750. torr) = 15 torr 2

The desired total pressure of 750 torr can also be expressed as 0.987 atm. So (0.987 atm)(15.0 L) PV ntotal = — = —— = 0.595 mol RT (0.08206 L atm mol−1 K−1)(303 K) ni = Xintotal nN = (0.751)(0.595) = 0.447 mol 2

nO = (0.149)(0.595) = 8.87 × 10−2 mol 2

nH O = (0.080)(0.595) = 4.8 × 10−2 mol 2

nSO = (0.020)(0.595) = 1.2 × 10−2 mol 2

Check Your Understanding A mixture of SO2(g) and SO3(g) is to be prepared with a total pressure of 1.4 atm. If the mole fractions of the gases are 0.70 and 0.30, respectively, what are the partial pressures? If the mixture is to occupy 2.50 L at 27°C, what mass of each gas is needed?

5.5

Stoichiometry of Reactions Involving Gases

When we discussed quantitative aspects of chemical reactions in Chapter 4, we emphasized the importance of ratios of moles. The ideal gas law provides a relationship between the number of moles of a gas and some easily measurable properties: pressure, volume, and temperature. So when gases are involved in a chemical reaction, the ideal gas law often provides the best way to determine the number of moles. Using the ideal gas law in a stoichiometry problem really doesn’t involve any new ideas. It just combines two kinds of calculations that you’ve already been doing. We’ll still do the stoichiometric calculation in terms of mole ratios, as always, and we’ll use the gas law to connect the number of moles of a gas with its temperature, pressure, and volume.


5.6

When an experiment required a source of carbon dioxide, a student combined 1.4 g of sodium bicarbonate (NaHCO3) with excess hydrochloric acid. If the CO2 produced is collected at 722 torr and 17°C, what volume will the gas occupy?

Strategy We are asked to find the volume of a gas, and we are given its pressure and temperature. We’ll assume that the gas behaves ideally. So if we knew the number of moles, we could easily use the gas law to get the volume we need. Looking a little closer, we should recognize this as a reaction stoichiometry problem because it asks us how much CO2 will be produced. The new wrinkle here is that it asks us to express the answer as a volume rather than as a mass or a number of moles. So we will first do a stoichiometric calculation to find the number of moles of CO2 produced and then

5.5 Stoichiometry of Reactions Involving Gases

139

use the gas law to find the volume of that amount of gas at the indicated temperature and pressure. As in any stoichiometry problem, we’ll need to start with a balanced equation for the reaction to be sure we use the correct mole ratio.

Solution NaHCO3 + HCl : NaCl + H2O + CO2 1 mol NaHCO3 1 mol CO2 1.4 g NaHCO3 × — × — = 0.017 mol CO2 84.0 g NaHCO3 1 mol NaHCO3 Now use the ideal gas law to solve for the volume. As always, we’ll need to remember to convert the temperature from °C to K and to be careful of pressure and volume units. (0.017 mol)(0.08206 L atm mol−1 K−1)(290 K) nRT = ——— V=— = 0.42 L = 420 mL (0.95 atm) P

Analyze Your Answer Our result of 420 mL is a relatively small volume—does this make sense? In this experiment, the reaction is similar to the one that occurs between baking soda and vinegar. Because 1.4 g of baking soda is not a large amount, it’s not surprising that the amount of gas generated is also small.

Check Your Understanding Suppose that 270 mL of CO2 at 1.03 atm and 25°C is produced by this reaction. Assuming that an excess of hydrochloric acid was present, how many grams of sodium bicarbonate were used?

STP Conditions From our discussion so far it should be obvious that to describe an amount of gas properly, we need to specify its temperature, volume, and pressure. If we want to compare various quantities of gas, it is convenient to use some standard set of conditions. To set a standard for an ideal gas, we might specify a temperature and a pressure; in that way, the volume would be directly proportional to the number of moles of gas. The conventional choices for the standard temperature and pressure of a gas, or STP, are a temperature of 0°C or 273.15 K and a pressure of 1 atm. Inserting these values into the ideal gas equation we find the molar volume (volume of one mole) of an ideal gas at STP. This standard molar volume provides a handy way to estimate gas properties or to help judge the reasonableness of a calculated value.

(1 mol)(0.08206 L atm mol−1 K−1)(273.15 K) nRT = ——— V=— = 22.41 L (1 atm) P This number provides a conversion factor for stoichiometric problems that include gases, provided that STP conditions are maintained. Example Problem 5.7 shows this type of problem.

E XAM PL E PRO BLEM

5 .7

One way to reduce air pollution is to remove potential pollutant gases from an exhaust stream before they are released into the air. Carbon dioxide can be removed from a stream of gas by reacting it with calcium oxide to form calcium carbonate. If we react 5.50 L of CO2 at STP with excess CaO, what mass of calcium carbonate will form?

Strategy You should recognize this as a reaction stoichiometry problem because it is asking us how much CaCO3 will be produced. As for any stoichiometry problem, we should write a balanced chemical equation for the reaction. Then convert the volume of gas into moles and proceed as usual. Because the gas volume given is at STP, we can use the molar volume we calculated above as a conversion factor.

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Chapter 5 Gases

Solution CO2(g) + CaO(s) : CaCO3(s) 100.1 g CaCO3 1 mol CaCO3 1 mol CO2 5.50 L CO2 × — × — × — = 24.6 g CaCO3 22.41 L CO2 1 mol CO2 1 mol CaCO3

Analyze Your Answer From the molar volume at STP, we can see that 5.5 L would be about one-quarter of a mole of CO2. So we should be making about onequarter mole of CaCO3, and our answer agrees with that.

Discussion Note that the conversion factor of 22.4 L = 1 mole is valid only when the gas in question is at STP. You should be very careful not to use that molar volume for any other conditions! We could also have worked this problem by using the ideal gas law to find the volume of the CO2, using the appropriate temperature and pressure for STP conditions. In this example, a gas is converted to a solid product rather than being emitted into the air. What implications might this have for engineering design? Can you think of situations in which this type of pollution control scheme might cause problems?

Check Your Understanding SO2 can be generated in the laboratory by heating ZnS with O2. (The other product is zinc oxide.) If 14.5 L of oxygen at STP is reacted with excess zinc sulfide, what volume of SO2 at STP is generated? At this point, we have seen how to carry out a range of calculations involving gases, all based on the ideal gas law. But though we have repeatedly said that we will assume the gas behaves ideally, we have not yet looked at what that statement implies in terms of molecular behavior. We have also not yet considered the many practical circumstances in which we might make large errors by assuming ideal gas behavior. We take up these challenges in the next section.

5.6


The ability to model the behavior of gases numerically was an important achievement that has broad applicability in many areas of science and engineering. The similarity of the behavior of different gases commonly encountered in the laboratory allowed the construction of models to describe how atoms must behave. But in many important practical settings, gases do not always behave ideally. So the simple ideal gas law calculations that we have been doing may not always be an accurate model of reality. Nonideal behavior is frequently seen in gases at very high pressures, such as might be found in the cylinder of an engine or in an industrial processing plant. Can an atomic perspective developed for ideal gases help explain these real world cases? Perhaps more importantly, can an improved model define quantitative methods for dealing with real gases such as those encountered in a running engine? In this section, we will develop a very useful model called the kinetic–molecular theory, which provides connections between the observed macroscopic properties of gases, the gas law equation, and the behavior of gas molecules on a microscopic scale. Once we understand this model, we can use it to see how gases at high pressure might deviate from ideal behavior.

The kinetic-molecular theory is also called the kinetic theory of gases or simply the kinetic theory.

Postulates of the Model Because observation of individual gas molecules is not generally feasible, we must begin with postulates that are accepted as reasonable when we construct a model.

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141

For the kinetic–molecular theory of gases, the following postulates form the bedrock of the model. 1. A gas is made up of a vast number of particles, and these particles are in constant random motion. 2. Particles in a gas are infinitely small; they occupy no volume. 3. Particles in a gas move in straight lines except when they collide with other molecules or with the walls of the container. Collisions with each other and with the walls of the container are elastic, so that the total kinetic energy of the particles is conserved. 4. Particles in a gas interact with each other only when collisions occur. 5. The average kinetic energy of the particles in a gas is proportional to the absolute temperature of the gas and does not depend on the identity of the gas. As we consider the implications of these postulates, we will also comment on their validity or the conditions under which they are not particularly accurate. Starting from these postulates, it is actually possible to derive the ideal gas law. Although that result is a great example of the synergy between scientific theory and experimental observations, it is a bit beyond our reach here. So instead we will consider the connections between the postulates of the kinetic theory and our practical experience with gases. How do these postulates help to explain the observed behavior of various gases? The statement that a gas consists of a large number of particles is consistent with what we know about the size of individual molecules. And the postulate that the molecules are in constant random motion is consistent with the observation that gases readily expand to fill their containers. The postulates are also consistent with what we said at the outset of this chapter about the origin of gas pressure. The molecules in a gas exert force when they collide with the wall of the container, and this accounts for the relationship between pressure and volume. As the volume of the container increases, the walls move farther apart and so each molecule will strike the wall less often. Thus, each molecule will exert less force on the wall in a given time interval, reducing the pressure. So the kinetic theory is consistent with the inverse relationship between P and V that we’ve seen in the gas law equation. How can the kinetic theory account for the effect of changes in temperature? The situation here is a little more complicated. We need to consider postulate #5 carefully: The kinetic energy of the particles in a gas is proportional to the absolute temperature of the gas and does not depend on the identity of the gas. We know from physics that kinetic energy is related to speed and to mass: KE = 1/2 mv2. So if the kinetic energy depends on the temperature, as stated in the postulate, then the speed of the molecules must be a function of temperature. More specifically, increasing the temperature must increase the speed of the molecules. When the molecules move faster, they will collide with the walls more often, and they will also impart a greater force on the wall during each collision. Both of these effects will lead to an increase in gas pressure with increasing temperature, as predicted by the gas law. If you were paying close attention to the statement of postulate #5, though, you may be wondering about the inclusion of the word “average” in “average kinetic energy.” That word is included because some molecules in a collection of randomly moving gas molecules move very fast, whereas others are virtually standing still at any given moment. A moment later the identity of fast moving and slow moving particles may have changed due to collisions, but there will still be both kinds. So it isn’t really correct to think about the molecules as having a particular speed, and we should actually say that the average molecular speed increases with increasing temperature. This situation is best approached using a distribution function to describe the range of speeds. Distribution functions are common mathematical devices used in science and also in economics, social sciences, and education. Your percentile scores on the SAT and other standardized tests are predicated on a distribution function for the students who take the exam. If you scored in the 85th percentile, for example, you would be expected to score higher than 85% of those taking the exam. This ranking is expected to hold independent 142

Chapter 5 Gases

Figure 5.6 ❚ The distributions of molecular speeds for CO2 molecules at three different temperatures are shown. The quantity on the y axis is the fraction of gas molecules moving at a particular speed. Notice that as the temperature increases, the fraction of molecules moving at higher speed increases.

300 K N(v)/N(total)

500 K

1000 K

0

500

1000 1500 Molecular speed (m/s)

2000

2500

of the number of students who have actually taken the test, so that your percentile is not revised every time another round of tests is administered to more students. The distribution function that describes the speeds of a collection of gas particles is known as the Maxwell-Boltzmann distribution of speeds. This function, which was originally derived from a detailed consideration of the postulates of the kinetic theory, predicts the fraction of the molecules in a gas that travel at a particular speed. It has since been verified by a variety of clever experiments that allow measuring the speed distribution in the laboratory. Let’s examine a plot of the Maxwell-Boltzmann distribution for a sample of CO2 gas at various temperatures, as shown in Figure 5.6. As with any unfamiliar graph, you should start by looking at what is plotted on each axis. Here the x axis shows a fairly familiar quantity: speed in meters per second. The y axis is a little less obvious. It shows the fraction of all the molecules in the sample that would be traveling at a particular speed under the given conditions. (The actual mathematical form of the distribution function is discussed in the MathConnections box.) We’ve already established that the distribution of speeds must be a function of temperature, and the figure bears this out. At high temperatures, we see that the peak in the distribution shifts toward higher speeds. This is entirely consistent with postulate #5, which states that the kinetic energy of the molecules must increase with temperature. Furthermore, postulate #5 also says that the kinetic energy does not depend on the identity of the gas. But the mass of the gas molecules clearly depends on their identity. For the kinetic energy to remain fixed as the mass changes, the speed distribution must also be a function of mass. Figure 5.7 illustrates this by showing the speed distribution for various gases at the same temperature. You can readily see that on average lighter molecules travel faster than heavier molecules.

Because we do not care in which direction the particles are going, our discussion here is based on speed rather than velocity. Speed does not depend on direction and is always positive.

Figure 5.7 ❚ Distributions of molecular speeds for four different gases at the same temperature (300 K). As in Figure 5.6, the quantity on the y axis is the fraction of gas molecules moving at a particular speed. Notice that as the molecular mass decreases, the fraction of the molecules moving at higher speed increases.

O2

N(v)/N(total)

N2 H2O

He

0

500


2000

5.6

2500


143

MathConnections The actual equation for the Maxwell-Boltzmann distribution of speeds is somewhat complicated, but it consists entirely of terms in forms that you should recognize from your math classes. N(v) Ntotal

3/2 2 −Mv /2RT — = 4p(M/2pRT ) v e 2

(5.11)

N(v) is the number of molecules moving with speeds between v and v + Δv, where Δv is some small increment of speed. Ntotal is the total number of molecules. So the left-hand side is the fraction of the molecules having speeds between v and v + Δv, which is also the quantity plotted on the y axis in Figures 5.6 and 5.7. M is the molar mass of the gas, R is the gas constant, and T is the temperature. (As in the gas law, this must be on an absolute scale.) Units for all of these quantities must be chosen consistently, with particular care taken to ensure that the exponent (–Mv 2/2RT ) has no units. Now look at Figure 5.6 or 5.7 again and notice that the independent variable on the x axis is the speed. So we want to think of the righthand side of the distribution function as an elaborate function of the speed, v. We see two terms containing v. The first is just v 2, which will obviously increase as v increases. The exponential term, though, varies as exp(–v 2), which will decrease as v increases. So the overall shape of the distribution function reflects a competition between these two terms. At small values of v, the v 2 term dominates, and the distribution function increases. At larger values of v, the exponential term takes over, and the overall function decreases. So, between those two limits, the distribution function goes through a maximum. The speed at which the peak in the plot occurs is known as the most probable speed, for fairly obvious reasons: more molecules travel at this speed than at any other. Although we will not do the derivation here, the most probable speed is given by Equation 5.12: vmp =

√

‾‾‾ 2RT

—

M

(5.12)

If you know some calculus you can do the derivation. Remember that the most probable speed is the maximum in the distribution function. To find this maximum, we would take the derivative of Equation 5.11 and set it equal to zero. We might also want to describe our distribution by talking about the average speed. Because the distribution function is not symmetric, the most probable speed and the average speed will not be the same. In particular, the existence of the “tail” on the distribution curve at high speeds will pull the average to a speed higher than the most probable value. The average speed actually turns out to be 1.128 times the most probable speed. Finally, sometimes we characterize the distribution by talking about something called the root-mean-square speed. This value is actually the square root of the average value of v 2. At first glance, it might sound like that should be the same as the average speed, but for the Boltzmann distribution, the root-meansquare speed is actually 1.085 times the average speed. (If this sounds wrong to you, try a simple test. Pick a few numbers at random and calculate their average. Then square all of your numbers and take the average of the squares. Finally, take the square root of that result and you’ll see that it is larger than the average of the original numbers.) The root-mean-square speed is useful if we are interested in kinetic energy because the average kinetic energy is given by 1 KEavg = — mv 2rms 2

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Chapter 5 Gases

(5.13)

Real Gases and Limitations of the Kinetic Theory In all of our calculations so far, we have begun by assuming that any gas or gas mixture behaves ideally. But the postulates of the kinetic theory—and the gas law itself—are strictly true only for purely hypothetical ideal gases. How good is the assumption that gases behave ideally, and under what kinds of conditions might we want to question this assumption? The answers to these questions can help us to develop further our understanding of gases at the molecular level. By learning how real gases deviate from ideal behavior, we will understand the nature of the kinetic–molecular theory better. We’ll start by taking a more critical look at some of the postulates of the theory. The assertion that gas molecules occupy no volume may seem at odds with reality because we know that all matter occupies space. What is actually implied by this postulate? The essential idea is that, compared to the volume of empty space between particles, the volume of the particles themselves is not significant. One way to assess the validity of this assertion is to define the mean free path of the particles. The mean free path is the average distance a particle travels between collisions with other particles. In air at room temperature and atmospheric pressure, the mean free path is about 70 nm, a value 200 times larger than the typical radius of a small molecule like N2 or O2. (For comparison, the molecules in a liquid typically have a mean free path roughly the same as a molecular radius.) When we consider the cubic relationship of volume to distance, the difference in volume is on the order of 2003 or 8 × 106. So the volume of empty space in a gas at room temperature and pressure is on the order of 1 million times greater than the volume of the individual molecules; the assumption in the kinetic theory that the volume of the molecules is negligible seems reasonable under these conditions. But under what circumstances would this assumption break down? We cannot generally increase the size of molecules substantially by changing conditions, so the only way the volumes become more comparable is to squeeze the molecules closer together by decreasing the volume. Thus under conditions of high pressure, when gases are highly compressed, the volume of individual molecules may become significant and the assumption that the gas will behave ideally can break down (see Figure 5.8a). This makes sense if we realize that one way to condense a gas into a liquid is to compress it, and of course we don’t expect the ideal gas law to hold for a liquid. The postulates also assert that gas molecules always move in straight lines and interact with each other only through perfectly elastic collisions. Another way to put this is to say that the molecules neither attract nor repel one another. If this were strictly true, then molecules could never stick together, and it would be impossible to condense a gas into a liquid or a solid. Clearly, there must be some attractive forces between molecules, a topic we will explore further in Chapter 8. The ideal gas model works because under many conditions these attractive forces are negligibly small compared to the kinetic energy of the molecules themselves. The fact that these forces Figure 5.8 ❚ The ideal gas model

(a) High pressure

(b)

Low temperature

5.6

breaks down at high pressures and low temperatures. (a) At high pressures, the average distance between molecules decreases, and the assumption that the volume of the molecules themselves is negligible becomes less valid. (b) At low temperatures, molecules move more slowly and the attractive forces between molecules can cause “sticky” collisions, as shown here by the paired molecules.


145

lead to the formation of liquids and solids points us to the conditions under which the forces become important. We know that gases will condense if we lower the temperature sufficiently. The reason for this is that at lower temperatures the kinetic energy of the molecules decreases, and eventually the strength of the attractive forces between molecules will become comparable to the kinetic energy. At that point, molecules will begin to experience sticky collisions, as shown in Figure 5.8b. This means that pairs or small clusters of molecules might stay in contact for a short time, rather than undergoing simple billiard ball collisions. This tendency for molecules to stick together, however fleetingly, will reduce the number of times that the molecules collide with the walls of the container, so the pressure of the gas will be lower than that predicted by the ideal gas equation.

Correcting the Ideal Gas Equation An empirical equation contains one or more adjustable parameters that are found from a best fit to observed data rather than from a theoretical model.

When the postulates of the kinetic theory are not valid, the observed gas will not obey the ideal gas equation. In many cases, including a variety of important engineering applications, gases need to be treated as nonideal, and empirical mathematical descriptions must be devised. There are many equations that may be used to describe the behavior of a real gas; the most commonly used is probably the van der Waals equation: qP

an2 +— r (V − nb) = nRT V2

(5.14)

Here, a and b are called van der Waals constants. Unlike the universal gas constant (R), the values of these van der Waals constants must be set for each specific gas. If a gas behaves ideally, both a and b are zero, and the van der Waals equation reverts to the ideal gas equation. Investigating the connection of this equation to the ideal gas equation further, we see that the term involving a accounts for the attraction between particles in the gas. The term involving b adjusts for the volume occupied by the gas particles. Table 5.2 provides the values of van der Waals constants for some common gases. Notice that larger molecules have larger values for b,

Table

❚

5.2

Van der Waals constants for several common gases

a (atm L2 mol−2)

b (L mol−1)

Ammonia, NH3

4.170

0.03707

Argon, Ar

1.345

0.03219

Carbon dioxide, CO2

3.592

0.04267

Helium, He

0.034

0.0237

Hydrogen, H2

0.2444

0.02661

Hydrogen fluoride, HF

9.433

0.0739

Methane, CH4

2.253

0.04278

Nitrogen, N2

1.390

0.03913

Oxygen, O2

1.360

0.03183

Sulfur dioxide, SO2

6.714

0.05636

Water, H2O

5.464

0.03049

Gas

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Chapter 5 Gases

consistent with its role in accounting for the volume of the molecules themselves. The largest value of a listed is that for HF. This means that compared to other gases, HF molecules interact with each other more and require larger corrections from ideal behavior. When we consider chemical bonding in the next few chapters, we will begin to understand why this is true. Example Problem 5.8 demonstrates the use of the van der Waals equation.


5.8

In Example Problem 5.1, we considered a CH4 storage tank with a volume of 49.0 L. When empty, the tank has a mass of 55.85 kg, and when filled, its mass is 62.07 kg. Calculate the pressure of CH4 in the tank at an ambient temperature of 21°C using both the ideal gas equation and the van der Waals equation. What is the percentage correction achieved by using the more realistic van der Waals equation?

Strategy We will use two different models to describe the same gas, and each model will be represented by its own equation. So we will do two independent calculations, solving for P in each case. We can find the mass of CH4 from the given data and then convert that mass into the number of moles by using the molar mass. And the van der Waals constants for CH4 can be found in Table 5.2. To calculate the percentage difference we will divide the difference of the two numbers by the value for the ideal gas case.

Solution Mass of CH4 = mass of full cylinder − mass of empty cylinder = 62.07 kg − 55.85 kg = 6.22 kg = 6220 g 1 mol CH4 Moles of gas = 6220 g CH4 × — = 388 mol CH4 16.04 g CH4 Ideal Gas: (388 mol)(0.08206 L atm mol−1 K−1)(294 K) nRT = ——— P=— = 191 atm 49.0 L V van der Waals: a = 2.253 L2 atm mol−2, b = 0.04278 L mol−1 n2 a nRT P=—−— V − nb V2 (388 mol)(0.08206 L atm mol−1 K−1)(294 K) = ——— 49.0 L − (388 mol)(0.0428 L mol−1) (388 mol)2(2.25 L2 atm mol−2) − —— (49.0 L)2 = 148 atm Percentage correction: 191 − 148 191

— × 100% = 23%

Discussion The correction here is quite significant. If you look at the calculated pressures, you should realize that they are quite high. Recall that the ideal gas law works best at low pressures and high temperatures. In this example, the pressure is sufficiently high that the ideal gas model is a rather poor description of the actual

5.6


147

behavior of the gas. Still, the simplicity of the ideal gas law makes it a good starting point for many calculations. In a practical application, it would be crucial for an engineer to recognize the conditions under which the assumption of ideal gas behavior might not be reasonable.

Check Your Understanding Under less extreme conditions, the corrections for nonideal behavior are smaller. Calculate the pressure of 0.500 mol of methane occupying 15.0 L at 25.7°C using both the ideal gas law and the van der Waals equation. Compare the percentage correction with that calculated in the example above.

INSIGHT INTO

5.7

Gas Sensors

The U.S. Environmental Protection Agency lists more than 180 substances as toxic air pollutants and specifies acceptable levels that can be less than 1 ppm in some cases. To track the effects of pollutants at such trace levels, first we must be able to measure their concentrations reliably. One way to describe the concentration of a gas, of course, is to report its pressure or partial pressure. Although a mercury barometer quite similar to one that might have been used hundreds of years ago will allow us to monitor the weather, it will not be of much help here. Instead we will need to rely on a number of more modern instruments that offer fundamental advantages over a barometer. A variety of devices now can allow us to measure exceedingly small pressures or changes in pressure. Some can also allow us to measure selectively the partial pressure of a particular component in a gas mixture. In this section, we will look at a couple of representative examples and try to show how their clever designs exploit the basic properties of gases to make incredibly sensitive measurements.

Capacitance Manometer The term manometer refers to any pressure-measuring device.

Figure 5.9 ❚ Schematic diagram of a capacitance manometer. The tube on the left is connected to the gas sample whose pressure is to be measured. As the gas molecules strike the diaphragm, the force of the collisions determines the space between the diaphragm and the fixed electrode assemblies.

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One sensor frequently used to measure pressure in the laboratory is the capacitance manometer. This type of pressure gauge, shown schematically in Figure 5.9, directly exploits the fact that pressure is force per unit area. A capacitor is an electronic component, as you may know from a physics class. One way to create a capacitor is to position two parallel metal plates with a space between them. In such an arrangement the distance between the plates determines the capacitance. A capacitance manometer is based on a capacitor in which one of these plates is actually a thin metal diaphragm and the other is fixed and rigid. The space Electrode assembly Diaphragm Baffle To electronics Gas sample

between the plates is kept under vacuum. We connect the sensor so that the other side of the diaphragm is exposed to the gas whose pressure we want to measure. The pressure of the gas produces a force on the diaphragm, causing it to bend toward the second (fixed) plate of the capacitor. This, in turn, changes the capacitance of the device. If we construct an electric circuit in which the capacitor is one component, we can measure a voltage in that circuit that will be related to the pressure applied to the diaphragm. To make the sensor easy to use, that circuit can be designed so that we can measure a single voltage whose value is directly proportional to the pressure. Monitoring the voltage allows us to see any changes in pressure. Capacitance manometers are most often used to measure pressures in the range of 0.001–1000 torr. This range includes moderate levels of vacuum frequently encountered in the laboratory. One advantage of this type of gauge is that it has virtually no moving parts; as long as the diaphragm is not damaged, the gauge should work reliably for very long periods of time.

Thermocouple Gauge Both the barometer and the capacitance manometer respond fairly directly to the force exerted by gas molecules. But not all pressure sensors work that way. A useful pressure-measuring device can be based on any gas property that varies predictably with pressure. Pressures between 0.01 and about 1.0 torr, for example, are frequently measured using a thermocouple vacuum gauge (Figure 5.10). A thin filament is heated by passing a fixed electric current through it, and the temperature of the filament is then measured using a device called a thermocouple. So how does this measurement tell us anything about pressure? The filament is in contact with the gas whose pressure we want to determine. As the molecules in that gas collide with the hot filament, they will tend to carry away some energy, reducing the filament temperature. The higher the pressure, the more molecules will collide with the filament, and the lower its temperature will be. The thermocouple produces a voltage that is related to the filament temperature and therefore to the gas pressure. A readout device connected to the gauge translates the voltage output into pressure units.

Ionization Gauge Capacitance manometers can be effective to pressures as low as about 0.001 torr. That is a very low pressure, just over a millionth of an atmosphere. But in certain types of experiments, such as studies of extremely clean metal surfaces, laboratory vacuums as low as 10−11 torr are necessary. At such low pressures, it is no longer feasible

Pressures on the order of 10 –14 torr are found in outer space.

Figure 5.10 ❚ Schematic diagram

Thermocouple junction

Filament Milliammeter

Millivoltmeter

of a thermocouple gauge. An electric current heats the filament, and a thermocouple monitors the filament’s temperature. Gas molecules collide with the filament and cool it. So the higher the pressure, the lower the filament temperature will be. A readout circuit usually converts the measured filament temperature into pressure units for display.

Power

5.7

Gas Sensors

149

Fine wire ion collector Cylindrical electron collection grid

Current amplifier Gas ion

Gas

+ Gauge display

Hot filament electron emitter

Figure 5.11 ❚ Schematic diagram of an ionization gauge. Electrons emitted from the hot filament collide with gas molecules and knock an electron free. The resulting cations are collected at the center of the gauge, and the current they produce is a measure of the gas pressure.

to measure the force produced by molecular collisions, so other types of gauges are needed. The most common of these is the ionization gauge, shown in Figure 5.11. A fixed electric current flows through the filament of the ionization gauge, heating it until it glows red. At these high temperatures, the metal atoms in the filament emit electrons. Some of these electrons then collide with gas molecules, and the resulting impact can knock an electron off the molecule, producing a positive ion. Suppose some or all of the gas present is N2. Electrons from the filament would then produce N2+ ions, through a process that we could write as e− + N2 : N2+ + 2 e− The resulting positive ions are then collected by a wire. Measuring the current flowing into this wire will measure how many ions were collected. The number of ions formed will be proportional to the number of gas molecules in the region of the gauge: more gas will allow more ions to be formed. Finally, we simply have to realize that the number of gas molecules present is proportional to the pressure. So the current, which is fairly easy to measure, will be proportional to pressure. In practice, the same electronics that read the current will also be calibrated so that they can display the reading directly as pressure.

Lawrence S. Brown

The photo shows (from left to right) a thermocouple gauge, a capacitance manometer, and an ionization gauge. The electronics needed to operate the gauges are not shown.

150

Chapter 5 Gases

Mass Spectrometer Both the capacitance manometer and the ionization gauge measure the total pressure of a gas. But if we want to measure the level of a pollutant in air, for example, we will really want to know its partial pressure. One way to measure partial pressures is with a mass spectrometer, which we saw earlier in Figure 2.4 (page 35). As in the ionization gauge, gas molecules are converted to positive ions by electron bombardment. But before the ions are collected, they are sent through a magnetic field that serves as a mass filter and sorts the stream of ions by mass. In this arrangement, the current that is measured consists only of ions with a chosen mass. This lets us monitor the partial pressure of a gas by looking at the ions of the corresponding mass. Control electronics allow these instruments to scan continuously over a range of masses and thus can tell us the partial pressure of a number of components in the same gas mixture.

FO C U S ON


Question A balloon is to be used to float instruments for observation into the upper atmosphere. At sea level, this balloon occupies a volume of 95 L at a temperature of 20°C. If helium gas is used to fill the balloon, how would you determine the maximum height that this balloon and its instruments will attain?

Strategy This is the type of question that the engineers who design apparatus such as this must answer. The law of nature here is that objects that are less dense float on those that are denser. The balloon will float upward until it reaches a height where its density is equal to that of the atmosphere. With this understanding, we can begin to define the variables we need to understand in this problem and note the relationships between those variables that would allow us to estimate an answer to such a problem.

Solution First identify what is known about the variables in the problem. In this case we have a gas that will be at low pressures, so we can assume it will behave as an ideal gas. We note that we will release this balloon from sea level where the pressure can be estimated as 1 atm. Therefore, we can establish the initial pressure, volume, and temperature. We also know the molar mass of helium. The key relationship is provided by the ideal gas law, which allows us to calculate density. To see this, think about n/V as the number of moles per unit volume. This is equal to the mass density divided by the molar mass. Replacing n/V with that relationship gives us a gas law in terms of density: rRT MM?P P = —, or r = — MM RT Looking at this set of equations, the importance of temperature and pressure is apparent. Because the key condition is the point at which the density of the helium in the balloon is equal to that of the atmosphere we must look up or determine the density of the atmosphere as a function of altitude. The source of this information might also provide the temperature and pressure of the atmosphere as a function of altitude. Knowing T and P provides the density of helium at various altitudes via the equation above. Having looked up the density of the atmosphere as a function of altitude, we must simply find the altitude at which this calculation equals the looked up value to estimate the maximum height that can be achieved by the balloon. We must recognize, however, that the mass of the instruments must be considered if we wish an accurate answer to this type of problem.

5.7

Gas Sensors

151

S U MMA RY The study of gases was extremely important in the development of a molecular description of matter, and still offers us the most direct glimpse into the connections between molecular behavior and macroscopically observable properties. In our everyday experience, no gas is more important than air. For those who live in urban areas, the reality of pollution plays a critical role in that experience. To understand gases, we need to think in terms of four key variables: pressure, temperature, volume, and quantity. Pressure can be measured over wide ranges by various devices, ranging from traditional barometers to modern mass spectrometers. To understand the origin of gas pressure, we must realize that the molecules in the gas exert a force when they collide with objects. Although any individual molecular collision imparts only a tiny force, the vast number of such collisions in a macroscopic sample of gas is enough to sum to a measurable value. For most applications, the ideal gas law, PV = nRT, provides the machinery necessary for quantitative understanding of gases. This equation includes the universal gas constant, R, in addition to the four variables we just noted. There are a number of ways

in which this equation can be used, but it is always important to watch the units used carefully. The equation itself is only valid if temperature is expressed on an absolute scale, usually the Kelvin scale. Although pressure and volume can be expressed in any convenient units, we must express R in units that are compatible with the other variables. The ideal gas law can be used to determine changes in a variable for a single gas—essentially recreating historic gas laws such as Boyle’s law or Charles’s law. It also allows us to extend our methodology for stoichiometry problems to include reactions involving gases. By applying the ideal gas law to each component of mixtures of gases, we can understand the concept of partial pressure and calculate the contribution of each gas in the mixture. The postulates of the kinetic–molecular theory provide a molecular scale explanation of what it means for a gas to be ideal. A careful examination of those postulates also points us to the conditions where gases are most likely to deviate from ideal behavior: high pressures and low temperatures. Under these conditions, empirical equations, such as the van der Waals equation, can account for observed deviations from ideal behavior.

KEY TERMS absolute temperature (5.6)

kinetic–molecular theory (5.6)

pressure (5.2)

atmosphere, atm (5.2)

manometer (5.7)

primary standard (5.1)

average speed (5.6)

Maxwell-Boltzmann distribution (5.6)

root-mean-square speed (5.6)

Avogadro’s law (5.3)

mean free path (5.6)

secondary standard (5.1)

barometer (5.2)

mole fraction (5.4)

Boyle’s law (5.3)

most probable speed (5.6)

standard temperature and pressure (STP) (5.5)

Charles’s law (5.3)

nonattainment area (5.1)

thermocouple gauge (5.7)

criteria pollutant (5.1)

part per million (5.1)

torr (5.2)

Dalton’s law of partial pressures (5.4)

partial pressure (5.4)

universal gas constant, R (5.1)

distribution function (5.6)

pascal, Pa (5.2)

van der Waals equation (5.6)

ionization gauge (5.7)

photochemical reaction (5.1)

volatile organic chemical (VOC) (5.1)



INSIGHT INTO Air Pollution 5.1 List two types of chemical compounds that must be present in the air for photochemical smog to form. What are the most common sources of these compounds? 5.2 When ozone levels in urban areas reach unhealthy levels, residents are typically urged to avoid refueling their cars during daylight hours. Explain how this might help to reduce smog formation. 5.3 In the production of urban air pollution shown in Figure 5.1, why does the concentration of NO decrease during daylight hours?

152

Chapter 5 Gases

5.4 VOCs are not criteria pollutants. Why are they monitored to understand urban air pollution? 5.5 Asphalt is composed of a mixture of organic chemicals. Does an asphalt parking lot contribute directly to the formation of photochemical smog? Explain your answer. 5.6 Do a web search to identify which (if any) of the EPA’s criteria pollutants are prevalent in your city or state. 5.7 One observable property of gases is the variability of density based on conditions. Use this observation to explain why hot air balloons rise.

■


Pressure 5.8 If you are pounding a nail and miss it, the hammer may dent the wood, but it does not penetrate it. Use the concept of pressure to explain this observation. 5.9 How do gases exert atmospheric pressure? 5.10 Why do mountain climbers need to wear breathing apparatus at the tops of high mountains such as Denali in Alaska? 5.11 If you had a liquid whose density was half that of mercury, how tall would you need to build a barometer to measure atmospheric pressure in a location where the record pressure recorded was 750 mm Hg?

5.26 A balloon filled to its maximum capacity on a chilly night at a carnival pops without being touched when it is brought inside. Explain this event. 5.27 A gas bubble forms inside a vat containing a hot liquid. If the bubble is originally at 68°C and a pressure of 1.6 atm with a volume of 5.8 mL, what will its volume be if the pressure drops to 1.2 atm and the temperature drops to 31°C? 5.28

■A

5.29

3 ■ A balloon filled with helium has a volume of 1.28 × 10 L at sea level where the pressure is 0.998 atm and the temperature is 31°C. The balloon is taken to the top of a mountain where the pressure is 0.753 atm and the temperature is –25°C. What is the volume of the balloon at the top of the mountain?

5.12 Water has a density that is 13.6 times less than that of mercury. How high would a column of water need to be to measure a pressure of 1 atm? 5.13 Water has a density that is 13.6 times less than that of mercury. If an undersea vessel descends to 1.5 km, how much pressure does the water exert in atm? 5.14 Does the vacuum above the mercury in the column of a barometer affect the reading it gives? Why or why not? 5.15

Gas pressures can be expressed in units of mm Hg, atm, torr, and kPa. Convert these pressure values. (a) 720. mm Hg to atm, (b) 1.25 atm to mm Hg, (c) 542 mm Hg to torr, (d) 740. mm Hg to kPa, (e) 700. kPa to atm ■

5.16 If the atmospheric pressure is 97.4 kPa, how much is it in mm Hg? In atm? 5.17 Why do your ears “pop” on occasion when you swim deep underwater? The Gas Law 5.18 When helium escapes from a balloon, the balloon’s volume decreases. Based on your intuition about stretching rubber, explain how this observation is consistent with the gas law. 5.19

■ A sample of CO2 gas has a pressure of 56.5 mm Hg in a 125-mL flask. The sample is transferred to a new flask, where it has a pressure of 62.3 mm Hg at the same temperature. What is the volume of the new flask?

5.20 A gas has an initial volume of 39 mL at an unknown pressure. If the same sample occupies 514 mL at 720 torr, what was the initial pressure? 5.21 When you buy a MylarTM balloon in the winter months in colder places, the shopkeeper will often tell you not to worry about it losing its shape when you take it home (outside) because it will return to shape once inside. What behavior of gases is responsible for this advice?

bicycle tire is inflated to a pressure of 3.74 atm at 15°C. If the tire is heated to 35°C, what is the pressure in the tire? Assume the tire volume doesn’t change.

5.30 How many moles of an ideal gas are there if the volume of the gas is 158 L at 14°C and a pressure of 89 kPa? 5.31

■ A newly discovered gas has a density of 2.39 g/L at 23.0°C and 715 mm Hg. What is the molar mass of the gas?

5.32 Calculate the mass of each gas at STP. (a) 1.4 L of SO2, (b) 3.5 × 105 L of CO2 5.33 What are the densities of the following gases at STP? (a) CF2Cl2, (b) CO2, (c) HCl 5.34 A cylinder containing 15.0 L of helium gas at a pressure of 165 atm is to be used to fill party balloons. Each balloon must be filled to a volume of 2.0 L at a pressure of 1.1 atm. What is the maximum number of balloons that can be inflated? Assume that the gas in the cylinder is at the same temperature as the inflated balloons. (HINT: The “empty” cylinder will still contain helium at 1.1 atm.) 5.35

A cylinder is filled with toxic COS gas to a pressure of 800.0 torr at 24°C. According to the manufacturer’s specifications, the cylinder may rupture if the pressure exceeds 35 psi (pounds per square inch; 1 atm = 14.7 psi). What is the maximum temperature to which the cylinder could be heated without exceeding this pressure rating?

5.36

■ Cylinders of compressed gases are often labeled to show how many “SCF” or “standard cubic feet” of gas they contain. 1 SCF of gas occupies a volume of 1 ft3 at a standard temperature and pressure of 0°C and 1 atm. A particular cylinder weighs 122 lbs when empty and 155 lbs when filled with krypton gas at 26°C. How many SCF of Kr does this cylinder contain?

■

Partial Pressure

5.22 Why should temperature always be converted to kelvins when working with gases in problem solving?

5.37 Define the term partial pressure.

5.23 What evidence gave rise to the establishment of the absolute temperature scale?

5.39 How does the mole fraction relate to the partial pressure?

5.24

■

Calculate the missing variable in each set. V1 = 2.0 L, T1 = 15°C, V2 = ?, T2 = 34°C V1 = ?, T1 = 149°C, V2 = 310 mL, T2 = 54°C V1 = 150 L, T1 = 180 K, V2 = 57 L, T2 = ?

5.25 Why is it dangerous to store compressed gas cylinders in places that could become very hot?

■


5.38 Define the term mole fraction. 5.40

■ What is the total pressure exerted by a mixture of 1.50 g H2 and 5.00 g N2 in a 5.00-L vessel at 25°C?

5.41 What is the total pressure (in atm) of a 15.0-L container at 28.0°C that contains 3.5 g N2, 4.5 g O2, and 13.0 g Cl2? 5.42 For a gas sample whose total pressure is 740 torr, what are the partial pressures if the moles of gas present are 1.3 mol N2, 0.33 mol O2, and 0.061 mol Ar?


153

5.43 A sample containing only NO2 and SO2 has a total pressure of 120. torr. Measurements show that the partial pressure of NO2 is 43 torr. If the vessel has a volume of 800.0 mL and the temperature is 22.0°C, how many moles of each gas are present? 5.44 A sample of gas is made entirely of carbon dioxide and water, and there are 259 moles of CO2 and 513 moles of water. If the total pressure of the sample is 21 atm, what is the partial pressure of each gas? 5.45

A sample of a smokestack emission was collected into a 1.25-L tank at 752 mm Hg and analyzed. The analysis showed 92% CO2, 3.6% NO, 1.2% SO2, and 4.1% H2O by mass. What is the partial pressure exerted by each gas? ■

5.46 Air is often modeled as dry air, ignoring the water present. If the atmospheric pressure outside is 0.944 atm and the pressure associated with water is 0.039 atm, what are the partial pressures of nitrogen and oxygen—assuming the composition given in Table 5.1. 5.47 In an experiment, a mixture of gases occupies a volume of 3.00 L at a temperature of 22.5°C. The mixture contains 14.0 g of water, 11.5 g of oxygen, and 37.3 g of nitrogen. Calculate the total pressure and the partial pressure of each gas. 5.48

■ An experiment is being carried out to test the removal of sulfur dioxide from gases being released by a power plant. The initial sample, which contains only carbon dioxide and sulfur dioxide, occupies a volume of 35 L at a temperature of 41°C and a pressure of 715 torr. After all of the SO2 has been removed, the sample occupies a volume of 23.5 L at a temperature of 29°C and a pressure of 715 torr. Determine the partial pressures of both CO2 and SO2 in the initial sample.

5.55 What volume of oxygen at 24°C and 0.88 atm is needed to completely react via combustion with 45 g of methane gas? 5.56

■ If boron hydride, B4H10, is treated with pure oxygen, it burns to give B2O3 and H2O.

2 B4H10(s) + 11 O2(g) : 4 B2O3(s) + 10 H2O(g) If a 0.050-g sample of the boron hydride burns completely in O2, what will be the pressure of the gaseous water in a 4.25-L flask at 30.°C? 5.57 N2O5 is an unstable gas that decomposes according to the following reaction: 2 N2O5(g) : 4 NO2(g) + O2(g) What would be the total pressure of gases present if a 10.0-L container at 22.0°C begins with 0.400 atm of N2O5 and the gas completely decomposes? 5.58 One way to generate oxygen is to heat potassium chlorate, KClO3 (the other product is potassium chloride). If 386 mL of oxygen at 41°C and 97.8 kPa is generated by this reaction, what is the minimum mass of KClO3 used? 5.59 Ammonia is not the only possible fertilizer. Others include urea, which can be produced by the reaction, CO2(g) + 2 NH3(g) : CO(NH2)2(s) + H2O(g). A scientist has 75 g of dry ice to provide the carbon dioxide. If 4.50 L of ammonia at 15°C and a pressure of 1.4 atm is added, which reactant is limiting? What mass of urea will form?

5.49 Use the web to determine the range of partial pressures that oxygen sensors must measure in the exhaust manifold of automobile engines. Why is it important for an engineer to know the amount of oxygen in an exhaust stream? 5.50 Use the web to research the types of fuel to air mixtures present in regular consumer driven automobiles. Calculate the partial pressures of the gases present based on the information you find. Stoichiometry with Gases 5.51 HCl(g) reacts with ammonia gas, NH3(g), to form solid ammonium chloride. If a sample of ammonia occupying 250 mL at 21°C and a pressure of 140 torr is allowed to react with excess HCl, what mass of NH4Cl will form? 5.52 Hydrogen gas is generated when acids come into contact with certain metals. When excess hydrochloric acid reacts with 2.5 g of Zn (the metal product is Zn2+), what volume of hydrogen gas is collected at a pressure of 0.93 atm and a temperature of 22°C?

Urea, CO(NH2)2

5.60 Consider the following reaction. 6 NiO(s) + 4 ClF3(g) : 6 NiF2(s) + 2 Cl2(g) + 3 O2(g) What mass of NiO will react with a sample of ClF3 gas that has a pressure of 250 torr in a 2.5-L flask at 20°C? ■

5.62

■ Magnesium will burn in air to form both Mg 3N 2 and MgO. What mass of each product would be found if burning a 3.11-g sample of magnesium to completion produces a combined total of 5.09 g of the two products?

5.63

2 ZnS(s) + 3 O2(g) : 2 ZnO(s) + 2 SO2(g)

■ During a collision, automobile air bags are inflated by the N2 gas formed by the explosive decomposition of sodium azide, NaN3.

If 620 kg of ZnS is to be reacted, what volume of oxygen at 0.977 atm and 34.0°C is needed (at a minimum) to carry out this reaction?

What mass of sodium azide would be needed to inflate a 30.0-L bag to a pressure of 1.40 atm at 25°C?

5.53 If you need 400.0 mL of hydrogen gas to conduct an experiment at 20.5°C and 748 torr, how many grams of Zn should be reacted with excess HCl to obtain this much gas? 5.54 The first step in processing zinc metal from its ore, ZnS, is to react it with O2 according to the reaction,

154

What volume of hydrogen gas, in liters, is produced by the reaction of 3.43 g of iron metal with 40.0 mL of 2.43 M HCl? The gas is collected at 2.25 atm of pressure and 23°C. The other product is FeCl2.

5.61

Chapter 5 Gases

2 NaN3 : 2 Na + 3 N2

■


Automakers are always investigating alternative reactions for the generation of gas to inflate air bags, in part because the sodium produced in the decomposition of NaN3 (see Problem 5.63) presents safety concerns. One system that has been considered is the oxidation of graphite by strontium nitrate, Sr(NO3)2, as shown in the equation below:

■

5 C(s) + 2 Sr(NO3)2(s) : 2 SrO(s) + 2 N2(g) + 5 CO2(g) Suppose that a system is being designed using this reaction and the goal is to generate enough gas to inflate a bag to 61 L and 1.3 atm at 23°C. What is the minimum mass of graphite that should be used in the design of this system? As one step in its purification, nickel metal reacts with carbon monoxide to form a compound called nickel tetracarbonyl, Ni(CO)4, which is a gas at temperatures above about 316 K. A 2.00-L flask is filled with CO gas to a pressure of 748 torr at 350.0 K, and then 5.00 g of Ni is added. If the reaction described above occurs and goes to completion at constant temperature, what will the final pressure in the flask be?

5.65

■

5.66

■ Ammonium dinitramide (ADN), NH4N(NO2)2, is being considered as a possible replacement for aluminum chloride as the oxidizer in the solid fuel booster rockets used to launch the space shuttle. When detonated by a spark, ADN rapidly decomposes to produce a gaseous mixture of N2, O2, and H2O. (This is not a combustion reaction. The ADN is the only reactant.) The reaction releases a lot of heat, so the gases are initially formed at high temperature and pressure. The thrust of the rocket results mainly from the expansion of this gas mixture. Suppose a 2.3-kg sample of ADN is detonated and decomposes completely to give N2, O2, and H2O. If the resulting gas mixture expands until it reaches a temperature of 100.0°C and a pressure of 1.00 atm, what volume will it occupy? Is your answer consistent with the proposed use of ADN as a rocket fuel?

5.67

Clouds of hydrogen molecules have been detected deep in interstellar space. It is estimated that these clouds contain about 1 × 1010 hydrogen molecules per m3 and have a temperature of just 25 K. Using these data, find the approximate pressure in such a cloud. ■

Kinetic Theory and Real Gases 5.68 State the postulates of the kinetic theory of gases. 5.69 Under what conditions do the postulates of the kinetic theory break down? 5.70 Which postulates of the kinetic theory are the most likely to break down under the conditions of high pressure present inside volcanoes? 5.71

Place these gases in order of increasing average molecular speed at 25 °C: Kr, CH4, N2, CH2Cl2.

■

5.72 Use kinetic theory to explain what happens to the pressure exerted by a gas as its temperature is increased. 5.73 The figure below shows the distribution of speeds for two samples of N2 gas. One sample is at 300 K, and the other is at 1000 K. Which is which?

N(v)/N(total)

5.64

0

500


2000

2500

5.74 Why do heavier gases move more slowly than light gases at the same temperature? 5.75 Define the term mean free path. 5.76 There are two parameters in the van der Waals equation for real gases. Which parameter corrects for the volume of gas molecules? Use the equation to explain your answer. 5.77 Calculate the pressure of 15.0 g of methane gas in a 1.50-L vessel at 45.0°C using (a) the ideal gas law and (b) the van der Waals equation using constants in Table 5.2. 5.78

You want to store 165 g of CO2 gas in a 12.5-L tank at room temperature (25°C). Calculate the pressure the gas would have using (a) the ideal gas law and (b) the van der Waals equation. (For CO2, a = 3.59 atm L2/mol2 and b = 0.0427 L/mol.)

5.79

■ Consider a sample of N2 gas under conditions in which it obeys the ideal gas law exactly. Which of these statements are true?

■

(a) A sample of Ne(g) under the same conditions must obey the ideal gas law exactly. (b) The speed at which one particular N2 molecule is moving changes from time to time.

NASA

(c) Some N2 molecules are moving more slowly than some of the molecules in a sample of O2(g) under the same conditions.

Interstellar hydrogen clouds, as seen by the Hubble Space Telescope

■


(d) Some N2 molecules are moving more slowly than some of the molecules in a sample of Ne(g) under the same conditions. (e) When two N2 molecules collide, it is possible that both may be moving faster after the collision than they were before.


155

Pressure Sensors

5.89

5.80 Scientists routinely use several different types of pressure gauges, and sometimes a single vacuum chamber may be equipped with more than one type of gauge. Why do you suppose that this is so? 5.81 What is actually measured in an ionization gauge pressure sensor? How is this actual measurement related to pressure? 5.82 Use the web to find out what type of sensor is used to determine the oxygen pressure in the exhaust manifold of automotive fuel injection system engines. 5.83 Do a web search to learn about at least one other type of pressure sensor not described in the chapter. Describe how the gauge operates in a way that your roommate could understand.

2 Al(s) + 6 H+(aq) : Al3+(aq) + 3 H2(g) Zn(s) + 2 H+(aq) : Zn2+(aq) + H2(g) The hydrogen produced has a volume of 0.147 L at 25°C and 755 mm Hg. What is the percentage of zinc in the alloy? 5.90 A mixture of Ar (0.40 mol), O 2 (0.50 mol), and CH 4 (0.30 mol) exerts a pressure of 740 mm Hg. If the methane and oxygen are ignited and complete combustion occurs, what is the final pressure of Ar, CO2, H2O, and the remainder of the excess reactant? What is the total pressure of the system? 5.91 The complete combustion of octane can be used as a model for the burning of gasoline: 2 C8H18 + 25 O2 : 16 CO2 + 18 H2O Assuming that this equation provides a reasonable model of the actual combustion process, what volume of air at 1.0 atm and 25°C must be taken into an engine to burn 1 gallon of gasoline? (The partial pressure of oxygen in air is 0.21 atm, and the density of liquid octane is 0.70 g/mL.)

Additional Problems 5.84

■

You have four gas samples:

1. 1.0 L of H2 at STP 2. 1.0 L of Ar at STP 3. 1.0 L of H2 at 27°C and 760 mm Hg

5.92

4. 1.0 L of He at 0°C and 900 mm Hg (a) Which sample has the largest number of gas particles (atoms or molecules)? (b) Which sample contains the smallest number of particles? (c) Which sample represents the largest mass? 5.85 A gas mixture contains 10.0% CH4 and 90.0% Ar by moles. Find the density of this gas at 250.0°C and 2.50 atm. 5.86

5.87

5.88

156

Liquid oxygen for use as a rocket fuel can be produced by cooling dry air to −183°C, where the O2 condenses. How many liters of dry air at 25°C and 750 torr would need to be processed to produce 150 L of liquid O2 at −183°C? (The mole fraction of oxygen in dry air is 0.21, and the density of liquid oxygen is 1.14 g/mL.) ■

■ A number of compounds containing the heavier noble gases, and especially xenon, have been prepared. One of these is xenon hexafluoride (XeF6), which can be prepared by heating a mixture of xenon and fluorine gases. XeF6 is a white crystalline solid at room temperature and melts at about 325 K. A mixture of 0.0600 g of Xe and 0.0304 g of F2 is sealed into a 100.0-mL bulb. (The bulb contains no air or other gases.) The bulb is heated, and the reaction above goes to completion. Then the sealed bulb is cooled back to 20.0°C. What will be the final pressure in the bulb, expressed in torr?

You work in a semiconductor production plant that relies on several chlorofluorocarbons in its manufacturing process. One day, you find an unlabeled gas cylinder, and you are assigned to figure out what is in the tank. First you fill a 1.000-L flask with the gas. At a pressure of 250.0 torr and a temperature of 25.00°C, you determine that the mass of the gas in the flask is 2.2980 g. Then you send the flask to an outside lab for elemental analysis, and they report that the gas contains 14.05% C, 44.46% F, and 41.48% Cl by mass. Find the molecular formula of this gas.

■

Chapter 5 Gases

■ A 0.2500-g sample of an Al-Zn alloy reacts with HCl to form hydrogen gas:

Mining engineers often have to deal with gases when planning for the excavation of coal. Some of these gases, including methane, can be captured and used as fuel to support the mining operation. For a particular mine, 2.4 g of CH4 is present for every 100. g of coal that is extracted. If 45.6% of the methane can be captured and the daily production of the mine is 580 metric tons of coal, how many moles of methane could be obtained per day? ■

5.93 Some engineering designs call for the use of compressed air for underground work. If water containing iron(II) ions is present, oxygen in the compressed air may react according to the following net ionic equation. Fe2+ + H+ + O2 : Fe3+ + H2O (a) Write the balanced net ionic equation. Remember that the amounts of each substance and the charges must balance. (b) Assume all of the oxygen from 650 L of compressed air at 15°C and 6.5 atm is lost by this reaction. What mass of water would be produced? (The mole fraction of oxygen in air is about 0.21.) (c) What will be the final pressure after the loss of the oxygen? 5.94 You are designing a warning system for a semiconductor fabrication facility that must have fewer than 1000 oxygen molecules per cm3 to operate successfully. While researching possible design options, you learn about a new sensor that can detect partial pressures of oxygen in the range of 1–2 mbar. Could this sensor be used as the basis for your warning system? Explain your answer. 5.95

Homes in rural areas where natural gas service is not available often rely on propane to fuel kitchen ranges. The propane is stored as a liquid, and the gas to be burned is produced as the liquid evaporates. Suppose an architect has hired you to consult on the choice of a propane tank for such a new home. The propane gas consumed in 1.0 hour by a typical range burner at high power would occupy roughly 165 L at 25°C and 1.0 atm, and the range chosen by the client will have 6 burners. If the tank under consideration holds 500. gallons of liquid propane, what is

■

■


to know the partial pressure of the H2 you generate as precisely as possible. What will you need to account for in your calculation? What things would you need to measure to make accurate calculations? What information would you need to look up?

the minimum number of hours it would take for the range to consume an entire tankful of propane? The density of liquid propane is 0.5077 kg/L.

FOCUS ON PROBLEM SOLVING EXERCISES 5.96 Before the introduction of unleaded gasoline, the amount of lead in severe pollution episodes in air could be measured at levels of 2.85 μg/m3. Describe how you could determine the number of lead atoms in the active space of your lungs (the air that is actively exchanged during a breath) if the volume of that space is 120 cm3. What, if anything, would need to be looked up? 5.97 A soft drink can label indicates that the volume of the soda it contains is 12 oz or 355 mL. There is probably some empty space at the top of the can. Describe what you can measure and how that measurement allows you to determine the actual density of the soda.

Cumulative Problems 5.101 Table 5.1 provides the mole percentage composition of dry air. What is the mass percentage composition of air? 5.102 Use the composition of air from Table 5.1 to calculate the average density of air at 25°C and 1 atm pressure. ■

5.104

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5.98 The best way to generate small amounts of oxygen gas in a laboratory setting is by the decomposition of potassium permanganate, 2 KMnO4(s) : K2MnO4(s) + MnO2(s) + O2(g). Suppose that you are carrying out an experiment for which you need exactly 2.00 L of oxygen. Describe how you would determine the mass of potassium permanganate needed to obtain this amount of oxygen. What would you have to measure in your experiment in addition to this mass? 5.99 An ore sample with a mass of 670 kg contains 27.7% magnesium carbonate, MgCO3. If all of the magnesium carbonate in this ore sample is decomposed to form carbon dioxide, describe how to determine what volume of CO2 is evolved during the process. What would have to be measured to predict the needed volume in advance? 5.100 Suppose that you are going to collect H2 gas produced from the reaction of an acid with a metal such as zinc. This reaction gives off heat, so the temperature of the solution will increase as the reaction proceeds. You need

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Consider a room that is 14 ft × 20 ft with an 8-ft ceiling. (a) How many molecules of air are present in this room at 20°C and 750 torr? (b) If a pollutant is present at 2.3 ppm, how many pollutant molecules are in this room?

5.103

Nitric acid can be prepared by bubbling dinitrogen pentoxide into water. N2O5(g) + H2O(,) : 2 H+(aq) + 2 NO3–(aq) (a) How many moles of H+ are obtained when 1.50 L of N2O5 at 25°C and 1.00 atm pressure is bubbled into water? (b) The solution obtained in (a) after reaction is complete has a volume of 437 mL. What is the molarity of the nitric acid obtained?

5.105

■ A 0.0125-g sample of a gas with an empirical formula of CHF2 is placed in a 165-mL flask. It has a pressure of 13.7 mm Hg at 22.5°C. What is the molecular formula of the compound?

5.106

■ Analysis of a gaseous chlorofluorocarbon, CClxFy, shows that it contains 11.79% C and 69.57% Cl. In another experiment, you find that 0.107 g of the compound fills a 458-mL flask at 25°C with a pressure of 21.3 mm Hg. What is the molecular formula of the compound?


157

6 OUTLINE 6.1

The Periodic Table and Atomic Structure

INSIGHT INTO

Incandescent and Fluorescent Lights 6.2

The Electromagnetic Spectrum

6.3

Atomic Spectra

6.4

Model of the Atom

6.5

The Pauli Exclusion Principle and Electron Configurations

6.6

The Periodic Table and Electron Configurations

6.7

Periodic Trends in Atomic Properties

6.8

INSIGHT INTO Modern Light Sources: LEDs and Lasers

This image shows the earth at night, as seen from orbiting satellites. The concentration of lights in industrialized nations, and even in specific cities, is very easy to see. The image is a composite of a large number of individual satellite photos. C. Mayhew and R. Simmon (NASA/GSFC), NOAA/ NGDC, DMSP Digital Archive

F

rom the fuels that power our planes and cars to the generation and control of air pollution, chemical reactions are central to our existence. What drives these reactions that are so much a part of who we are and what we do? How can we predict which reactions are likely to happen? To answer questions such as these, we will need an understanding of the chemical bonds that link atoms into molecules. To understand the bonds between atoms, first we must refine our descriptions of the atoms themselves. Fortunately, there are tools that allow us to observe atoms and other tools to help us organize what we have observed and learned. No tool is more important to building our understanding of chemistry than the periodic table. Its structure is closely related to the nature of atoms. Our understanding of atomic structure, in turn, has influenced the development of countless modern technologies. Nowhere is this influence more direct than in the lighting industry.

Chapter Objectives


158

After mastering this chapter, you should be able to ❚ describe similarities and differences between various light sources. ❚ describe waves in terms of frequency, wavelength, and amplitude. ❚ interconvert between the frequency and wavelength of light.

❚ relate the frequency, wavelength, and amplitude of light to characteristics such as color and brightness. ❚ describe the photoelectric effect by stating what sort of experiment is involved and what results are seen. ❚ explain how the results of the photoelectric effect experiment are consistent with a photon model of light. ❚ use Planck’s equation to calculate the energy of a photon from its wavelength or frequency. ❚ use ideas about conservation of energy to explain how the observation of atomic spectra implies that electrons in atoms have quantized energies. ❚ use an energy-level diagram to predict the wavelengths or frequencies of light that an atom will absorb or emit, or use the observed wavelengths or frequencies to determine the allowed energy levels. ❚ describe similarities and differences between the Bohr model and the quantum mechanical model of atomic structure. ❚ recognize how quantum numbers arise as a consequence of the wave model. ❚ define the term orbital. ❚ identify an orbital (as 1s, 3p, etc.) from its quantum numbers, or vice versa. ❚ list the number of orbitals of each type (1s, 3p, etc.) in an atom. ❚ sketch the shapes of s and p orbitals and recognize orbitals by their shapes. ❚ rank various orbitals in terms of size and energy. ❚ use the Pauli exclusion principle and Hund’s rule to write electron configurations for atoms and ions of main group elements. ❚ explain the connection between valence electron configurations and the periodic table. ❚ define the following properties of atoms: atomic radius, ionization energy, and electron affinity. ❚ state how the above properties vary with position in the periodic table.

INSIGHT INTO

6.1

Incandescent and Fluorescent Lights

Electric lights are one of the hallmarks of modern civilization. Few visit Las Vegas or Times Square without being awed by the lights. As we set out to explore how those lights operate, we’ll choose a rather unlikely point of departure: a blacksmith or metalworker heating a piece of iron in a furnace. You may have seen someone playing the part of a 19th century blacksmith in a historic village or open-air museum. As the blacksmith works a piece of iron, he begins with a cold, dull gray piece of metal. After a short time in the hot fire, the iron slowly becomes a dull red color. After further heating, the metal shines bright orange, and finally it appears almost white. Why does the metal change from dull gray to orange as it is heated? This question may seem to have little to do with a lightbulb until we consider the material that makes up the bulb’s filament. It, too, is a metal, usually tungsten. Thus the heating of iron in a blacksmith’s furnace provides some clues to the technology behind some types of light bulbs. Every light bulb operates on the same principle. Electricity is supplied to “excite” a substance, and that excitation somehow gives rise to light. But what exactly does it mean to “excite” a substance? This is one of the questions that we will discuss in this chapter, and the answer will lead us into a consideration of the nature of the atom itself. Just what happens in the excitation process depends on the type of light bulb. The filament in a standard incandescent light bulb is a length of very thin tungsten wire. When an electric current is passed through the filament, the resistance of the wire

The filament in a typical 60-watt bulb is about 0.01 inch thick. It would be about 6.5 feet long if uncoiled. When the bulb is lit, the filament temperature reaches about 4000°F.

6.1 Incandescent and Fluorescent Lights

159

Figure 6.1 ❚ The diagram shows the important components of a standard incandescent light bulb and lists the chemical substances from which many of the components are made.

Glass

Bulb Filament

Tungsten

Supports

Molybdenum or tungsten

Button Button rod Stem

Copper

Press Tube

Copper Nickel Iron

Lead wire

Cement

Exhaust tube Base

Shell contact Glass insulator Center contact

Brass Solder

Compact fluorescent bulbs operate in the same way as traditional fluorescent tubes. The main difference is that the compact fluorescent bulb is engineered to be compatible with standard lamps or fixtures that were initially designed for incandescent bulbs.

Figure 6.2 ❚ The diagram shows the important components and processes in a standard fluorescent light.

Soda Lead Potash Nitre Arsenic Silica Calcium Cobalt Feldspar Manganese

Malachite green Marble Bakelite Alcohol Resin Zinc Copper

Lead Tin

causes some energy to be dissipated as heat. The tungsten is heated until it is “white hot,” and it emits visible light (Figure 6.1). Because the tungsten wire is so thin, it heats up quickly to a high temperature. But it is a metal that is behaving similarly to the blacksmith’s iron. The specific color of light emitted by either metal is related to its temperature. The dull red color of the iron as it is first heated suggests that, although warm, the piece is not yet especially hot. The move toward orange as the piece remains in the fire, or the white light from the tungsten filament suggests higher temperatures. With the two examples we have noted so far, the range of colors is not that wide, but the progression of color is similar to the colors of the rainbow that begin with red and move toward orange. Is there anything more than coincidence associated with this observation? We will investigate the nature of electromagnetic radiation to gain insight into this question. Fluorescent lights operate differently. Instead of heating a metal, an arc of electricity travels through a gas. An arc occurs whenever electricity passes through a medium that does not conduct current very well. The discharge of static electricity that sometimes occurs when you touch a light switch provides a painful, personal example of an electric arc. Figure 6.2 illustrates the processes in a fluorescent light bulb. In such a bulb, specific atoms such as mercury interact with the electricity and become “excited.” This means that they are in a state of higher than normal energy. To return to their normal state, these excited atoms emit light. That light interacts with a material, called a phosphor, that coats the inside surface of the fluorescent tube.

1 Electricity supplied to the poles builds a large voltage difference.

4 Phosphor absorbs UV light and emits visible light.

2 Electrical arc excites atoms (often Hg) in the tube. 3 Excited mercury atom emits UV (and other) radiation.

160

Chapter 6 The Periodic Table and Atomic Structure

This phosphor then emits light of many colors, giving rise to what we perceive as white light coming from the bulb. What happens when atoms get excited or relax? Why does the bulb need to use the phosphor? Could elements other than mercury be used? All of these questions can be answered as we also establish some of the foundation upon which to build a deeper understanding of chemical bonding. Therefore, we proceed to learn how we view atoms in modern chemistry and physics and how that understanding provides the insight needed to light the world. Because the interaction of light with atoms played a crucial role in the development of our model for the structure of atoms, we begin with a look at the properties of light itself.

6.2


Visible light is a more accurate term for what we usually refer to simply as light. The light that our eyes can detect comprises only a small portion of the possible electromagnetic spectrum and accounts for only a small part of the emission of most light bulbs. Other familiar forms of electromagnetic radiation include radio waves, microwaves, and X-rays. The origin of the word electromagnetic lies in the nature of light. Historically, light has been described as a wave traveling through space. One component is an electric field, and another is a magnetic field (Figure 6.3). To understand the wave nature of light, then, we should probably be familiar with the general features of waves.

The Wave Nature of Light Many of the features of light that we encounter can be explained as properties of waves. The central characteristics of a wave can be defined by four variables: wavelength, frequency, velocity, and amplitude. Figure 6.4 illustrates the definition of these terms. The wavelength of any wave is the distance between corresponding points on adjacent waves. In the illustration, for example, the wavelength (designated as l and measured in units of length) is defined as the distance between peaks. We could also define the wavelength as the distance between valleys, and the value would be the same. The amplitude is the size or height of the wave.

Figure 6.3 ❚ Light is one form of electromagnetic radiation, which consists of oscillating electric and magnetic fields.

Amplitude Wavelength, l Electric vector Magnetic vector Nodes

Direction of propagation

6.2


161

Figure 6.4 ❚ Two sine waves are shown, and the meanings of some important terms are illustrated. Because light has a fixed speed, its frequency and wavelength are always inversely related. Amplitude is independent of frequency; the choice of which of these waves would have larger amplitude was made arbitrarily.

Wavelength λ

Amplitude

ν is the Greek letter nu.

Low frequency Long wavelength Small amplitude

High frequency Short wavelength Large amplitude

The lower wave in Figure 6.4 has a larger amplitude than the upper wave. The frequency is the number of complete cycles of the wave passing a given point per second. Frequency is usually designated as ν and measured in units of 1/s, or hertz (Hz). Frequency and wavelength are not independent of one another. The speed of light in a vacuum (designated by c and measured in units of distance/time) is a constant of nature that has been measured with impressive precision and is now defined to have an exact value: c = 2.99792458 × 108 m s−1 Because the speed of light is a constant, the number of waves that will pass a certain point in time will be inversely related to the wavelength. Long wavelengths will have fewer cycles pass by in one second, so they will have lower frequencies (Figure 6.4). This relationship is expressed mathematically in Equation 6.1: c=l×n

(6.1)

MathConnections If we want to describe the detailed shape of a wave, we need to write a mathematical equation for it. Let’s consider a very simple example in which the wave happens to follow the shape of a sine function. (This turns out to be true for many naturally occurring waves.) Imagine that we take a snapshot of the wave at a particular instant and picture it as an x – y graph. The wave originates at x = 0 and travels along the x axis. The y values represent the height of the wave at each point. We could write a very general equation for this wave: y = A sin (bx) A is clearly the amplitude of the wave, but what about b? Let’s assume that we want to evaluate the sine function in terms of radians. That means that the quantity bx must have no units. Because x is the distance from the origin, it would have units of length. So the b in our equation must have units of 1/length. If we know the wavelength of the wave—which we could read off the graph—then we can also fix the value of b. The wavelength of the sin (x) function is 2π. So, suppose that the wavelength of our wave is 450 nm, in the range of blue light. For our equation to match the wave, we require that b × 450 nm = 2p

162


or b = 2p/450 nm More generally, we can see that the b in our equation must always be 2p/l. If we wanted to describe a more complicated wave, we might have to use a combination of several sine (or cosine) functions, each of the same general form as we see here but with different wavelengths. By combining enough such functions, we could describe a waveform of arbitrary complexity. If we wanted to describe a wave as a function of time rather than position, we could proceed in a similar fashion.

Speed, wavelength, frequency, and amplitude are characteristics that can describe any wave. But because we are interested in light, it should help to compare these quantities with some familiar properties of light. If someone asks you to compare two different lights, one of the things you would probably consider is how bright each one appears. How does this relate to the wave parameters we’ve defined? In the wave model for light, the amplitude determines the brightness of the light: the larger the amplitude of the wave, the brighter the light will appear. Another property that you would surely consider in comparing two lights is color. In the wave model, both wavelength and frequency correspond to the color of the light. How can these two seemingly different properties of the wave both correspond to color? The answer can be seen in Equation 6.1. Because the speed of light must remain constant, specifying either the frequency or the wavelength automatically fixes the value of both of those properties of a light wave. So, we could say that the frequency and the wavelength are simply two different ways of expressing the same information. The wave nature of light also helps to explain many of the phenomena we experience with light. Refraction is the bending of a wave when it passes from one medium to another of different refractive index. One common manifestation of refraction is the rainbow, in which white light is separated into colors. In passing from air to water or glass, the speed of light changes, and the light bends at an angle that depends on its wavelength. Thus the visible spectrum is separated into its colors, and the colors are dispersed in order of their corresponding wavelengths. Figure 6.5 illustrates this phenomenon. Because of wave-like phenomena such as refraction, it is common to categorize electromagnetic radiation in terms of its wavelength or frequency. Figure 6.6 shows the entire spectrum, including information about both frequency and wavelength.

Mathematically, the intensity of a light wave is actually proportional to the square of its amplitude.

Increasing wavelength Slits to isolate thin beam White light

Prism

Figure 6.5 ❚ When white light passes through a prism, it is separated into its constituent colors by refraction.

6.2


163

Energy Increases 1024

1022

1020

g-rays 10⫺16

10⫺14

1018

1016

X-rays 10⫺12

10⫺10

1014

UV 10⫺8

1012 IR

10⫺6

1010

108

106

104

Microwave FM AM Radiowaves

10⫺4

10⫺2

100

102

102

100

n (Hz)

Long radio waves 104

106

108

l (m)

Visible spectrum

400

500 Energy increases

600 Wavelength increases

700

l (nm)

Figure 6.6 ❚ Visible light is just the small portion of the entire electromagnetic spectrum that our eyes can detect. Other regions of the spectrum are used in many technological applications, including remote controls, cell phones, and wireless networks.

In addition to visible light, the spectrum includes X-ray, ultraviolet (UV), infrared (IR), microwave, and radio wave radiation; the visible falls between UV and IR. This order lists the regions of the spectrum in increasing wavelength. We will often find that we can measure or are provided with the wavelength of light but need to express it as a frequency. Example 6.1 shows this type of conversion.

E XAM PL E PRO BLEM

6 .1

Neon lights emit an orange-red colored glow. This light has a wavelength of 670 nm. What is the frequency of this light?

Strategy We know that wavelength and frequency are related through Equation 6.1. In using this equation, we will need to be careful of units. Because the wavelength is given in nanometers, we will convert it to meters so that it is consistent with the meters per second in the speed of light.

Solution 1m 670 nm × —————— = 6.7 × 10−7 m 9 1 × 10 nm c We know that c = lν, so, rearranging, ν = ——. Then we simply substitute the values for l the speed of light and the wavelength in the equation: 2.998 × 108 m s−1 ν = ————————— = 4.5 × 1014 s−1 6.7 × 10−7 m

Analyze Your Answer It is difficult to have a physical intuition about something like the frequency of light. But with some practice we can develop an idea of the expected order of magnitude. First, notice that to find the frequency, we took a large number (the speed of light) and divided it by a small number (the wavelength). So we should expect a very large numerical answer. It is also handy to become familiar with the typical magnitudes of some quantities. A glance at Figure 6.6, for example, will show us that the frequency of visible light is on the order of 1014 s−1, so our answer

164


seems plausible. These kinds of checks can be very handy because the most common errors in problems like this involve mishandling of units, leading to results that can be off by several orders of magnitude.

Check Your Understanding When trace amounts of xenon gas are added to neon lights, the color becomes blue-green. If the wavelength of this light is 480 nm, what is the frequency?

Generally, we think of a lamp as something emitting visible light. But for some specific applications, bulbs are designed to produce radiation in different regions of the electromagnetic spectrum. Heat lamps in fast food restaurants, for example, emit infrared radiation to keep French fries warm. Compared with ordinary incandescent bulbs, heat lamps are designed to emit a range of wavelengths shifted toward the IR region. These bulbs appear red to our eyes because that is the visible color that is adjacent to the infrared. Moving to the opposite end of the visible spectrum, lamps designed to emit ultraviolet radiation are used both as “black lights,” causing various dyes to glow, and in tanning beds. Our eyes do not detect ultraviolet light, but these lamps appear purple because the range of wavelengths they emit overlaps the violet end of the visible spectrum. Microwave ovens, radios, televisions, wireless telephones, and wireless computer networks all function by exploiting other specific regions of the electromagnetic spectrum.

The Particulate Nature of Light The wave model of electromagnetic radiation can explain many observed properties of light. In the early 1900s, however, scientists developed more sophisticated apparatus and carried out experiments that challenged the wave model of light. One pivotal group of experiments involved the observation of what came to be known as the photoelectric effect, in which light strikes a piece of metal and causes electrons to be ejected. Figure 6.7 is a schematic diagram of a typical photoelectric effect 1 Photons of sufficient energy strike a metal surface and electrons are ejected.

Electron

Figure 6.7 ❚ In the photoelectric

2 The ejected electrons are collected by the anode.

effect, light strikes the surface of a metal and electrons are ejected, as shown in the inset. This is exploited in photoelectric cells: light strikes a surface called the photocathode, and a wire (the anode) collects the emitted electrons. (A positive voltage applied to the anode ensures that the electrons travel in the desired direction.) The resulting current can be measured easily and is related to the intensity of the incident light.

Electron (–) Photo cathode (–)

Light

Anode (+)

Meter (current) 3 The movement of the photoelectrons causes a current to flow, and allows the light intensity to be monitored.

6.2


165

experiment. At the heart of the experiment is light shining on a piece of metal. If the metal is enclosed in a container under vacuum so that there are very few surrounding gas molecules to interfere, then we can detect electrons emitted from the metal under certain conditions. (The metal itself is grounded so that it does not develop a large positive charge during this process.) As we learned in Section 2.2 (page 33), electrons carry a negative charge. This makes it fairly easy to measure how many electrons are ejected and even how fast they are going. So in studying the photoelectric effect, the experimenter could vary the color or intensity of the light and see how the number of electrons emitted or their kinetic energy might change. The experiment could also be repeated using different metals. How can this photoelectric effect be understood? The simplest explanation is that energy from the light is transferred to the electrons in the metal. If an electron picks up enough energy, it can break free from the surface of the metal. The more energy given to the electron, the faster it will travel after it leaves the metal. Figure 6.8 summarizes the sort of results obtained in these experiments. A close look at these graphs reveals that the detailed observations are not what we would expect based on the wave model for light. In a wave description of a light source, the amplitude of the wave or the intensity of the light determines the energy. So we would expect that as the light gets brighter, the kinetic energy of the outgoing electrons should increase. But Figure 6.8(d) shows that the kinetic energy of the electrons is actually independent of the intensity of the incoming light. Only the number of electrons detected increases when the intensity goes up (graph b). But the kinetic energy does depend on the frequency or color of the light used (graph c). This is also at odds with the wave model. So these experiments (and others) caused scientists to reexamine the nature of light. # e–

# e– υ > υ0

(a)

υ0

υ

KE

(b)

I

KE υ > υ0

(c)

υ0

υ

(d)

I

Figure 6.8 ❚ The graphs sketched here show the most important observations from experiments on the photoelectric effect. Below some threshold frequency (ν0), no photoelectrons are emitted. For frequencies greater than ν0, the number of electrons emitted is independent of frequency (graph a) but increases with increasing light intensity (graph b). The kinetic energy of the emitted electrons, on the other hand, increases linearly with frequency (graph c) but is independent of intensity (graph d).

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Despite elaborate efforts to make these observations of the photoelectric effect conform to existing wave concepts, consistent explanations of light based solely on waves ultimately failed. The only way to explain all of the experimental results was to invoke the notion of wave-particle duality, which says that, in some situations, light is best described as a wave, whereas in other cases, a particle description works better. It’s important to realize that this does not mean that there are two different kinds of light! It simply means that neither the wave model nor the particle model provides an accurate description of all of the properties of light. So in a given situation, we use whichever model works best to describe the properties we are interested in. Someday, advances in our understanding of light may produce a new model that reconciles the wave-particle conflict. But until that happens, we take a pragmatic approach and use the model that’s best suited to the problem at hand. The context in which we normally think of light as a particle is when it is imparting energy to another object. Albert Einstein proposed that light could be described as a collection of packets of energy, called photons. Bright light has many photons, whereas dim light has few. The energy of a photon of light has been shown to be proportional to its frequency, leading to a simple equation: E = hn

(6.2)

Here E represents the photon energy, and n is the frequency of the light. The h term is a constant, called Planck’s constant, after German scientist Max Planck. Its value is

Recall that in Chapter 3 (p. 68) we used hν to indicate that light is required in a photochemical reaction.

h = 6.626 × 10−34 J s. Because frequency is related to wavelength we can substitute for ν in this equation, using Equation 6.1: hc E = —— l

(6.3)

Thus, if we know either the wavelength or the frequency of radiation, we can determine the energy of a photon of that radiation, as shown in Example Problem 6.2.

E X AM PL E PROB L E M

6.2

The laser in a standard laser printer emits light with a wavelength of 780.0 nm. What is the energy of a photon of this light?

Strategy We know the connection between photon energy and wavelength, which is given by Equation 6.3. Again, care with units requires conversion from nanometers to meters.

Solution 1m 780.0 nm × —————— = 7.800 × 10−7 m 1 × 109 nm 6.626 × 10−34 J s × 2.998 × 108 m s−1 E = ————————————————— = 2.547 × 10−19 J 7.800 10−7 m

Analyze Your Answer The result is a very small number. But we should realize that Planck’s constant is incredibly small—many orders of magnitude smaller than the other quantities in the problem. So a very small energy is likely. Beyond that, we can again rely on a general sense of the magnitude of the quantity we are calculating. For visible light, photon energies are typically on the order of 10−19 J, which makes our answer seem plausible.

6.2


167

Check Your Understanding An infrared laser for use in a fiber-optic communications network emits at a wavelength of 1.2 mm. What is the energy of one photon of this radiation?

Be careful not to confuse the photon energy with the intensity or brightness of the light. If you remember that photon energy is another way of specifying the color of the light, you will see that it is independent of intensity.

The idea that light is made up of a collection of particles, or “chunks of energy,” does not seem consistent with our everyday experience. To understand why, look at the magnitude of the photon energy calculated in Example Problem 6.2: 10−19 J. Compare that with the amount of energy emitted by an ordinary lamp. Even a modest 60-watt bulb emits 60 J of energy each second. So the lights that you are familiar with all consist of vast numbers of photons. As the light gets brighter, the number of photons emitted per second increases. But because each photon is such a small additional amount of energy, it appears to us as if the energy coming from the lamp varies continuously rather than in discrete increments or steps. Though not readily discernible to us, the energy of a single photon can be very significant from an atomic or molecular vantage point. Atoms and molecules are very small, after all, and their typical energies are small as well. Consider the results presented earlier for the photoelectric effect experiment. How can the concept of photons help us to explain those observations? Imagine that we could do the experiment by shooting a single photon at the metal. That means we would be supplying an amount of energy (equal to hν) to the metal. If this energy is enough to overcome the force that the metal atoms exert to hold onto their electrons, then when the metal absorbs the photon, it will be able to eject one electron. If the energy of the photon is smaller than the binding energy holding the electrons to the metal, then no electron can be emitted. The fact that there is a threshold frequency (ν0 in Figure 6.8), below which no electrons are detected, is evidence that the photoelectric process is carried out by individual photons. Below the threshold frequency, a photon does not have enough energy to overcome the binding energy and eject an electron. Why does the electron kinetic energy depend on the frequency of the light? This is really a question of conservation of energy. For a particular metal, the electron binding energy is fixed. For any frequency above ν0, a portion of the energy of the absorbed photon will be used to overcome the binding energy. Then, whatever energy remains will be transferred to the electron as kinetic energy. So, as the frequency increases, the photon energy increases. This leads to an increase in the observed kinetic energy of the ejected electrons.

E XAM PL E PRO BLEM

6 .3

In a photoelectric effect experiment, ultraviolet light with a wavelength of 337 nm was directed at the surface of a piece of potassium metal. The kinetic energy of the ejected electrons was measured as 2.30 × 10−19 J. What is the electron binding energy for potassium?

Strategy We know that energy must be conserved, and we can use this idea to set up an equation relating the photon energy to the kinetic energy and the binding energy. Because the photon is absorbed in this process, we must be able to account for its energy. Some of that energy does work to overcome the electron binding energy, and the rest appears as the kinetic energy of the ejected electron. We can write this as a simple equation: E photon = Binding E + Kinetic E We know how to find the photon energy from the wavelength, and we know the kinetic energy, so we can solve for the binding energy. As always, we must be careful with the units on all quantities involved.

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Solution First find the photon energy. 1 m = 3.37 × 10−7 m 337 nm × ———— 109 nm (6.626 × 10−34 J s)(2.998 × 108 m s−1) hc Ephoton = —— = ————————————————— = 5.89 × 10−19 J l 3.37 × 10−7 m Then use the conservation of energy relationship to find the desired binding energy: Ephoton = Binding E + Kinetic E So, Binding E = Ephoton − Kinetic E = 5.89 × 10−19 J − 2.30 × 10−19 J = 3.59 × 10−19 J

Analyze Your Answer The wavelength used here is slightly shorter than those for visible light, so photon energy in the mid to upper 10−19 J range seems reasonable. The binding energy we found is smaller than the photon energy, as it must be. So it appears likely that we have done the calculation correctly. Check Your Understanding The electron binding energy for chromium metal is 7.21 × 10−19 J. Find the maximum kinetic energy at which electrons can be ejected from chromium in a photoelectric effect experiment using light at 266 nm. By combining the concepts we’ve just developed, we can begin to consider some of the issues about light bulbs raised in the opening section. We noted that when a blacksmith at a museum heats iron, it first glows red and then turns orange. Looking at the spectrum we see that red light has a longer wavelength than does orange light. So the energy of a photon of orange light is greater than that of a photon of red light. The iron changes from red to orange because the heating process has increased its energy. An object that emits orange light is at a higher temperature than one emitting red light. Now that we’ve discussed the nature of light itself, we need to consider the way that light allows us to investigate atoms.

6.3

Atomic Spectra

You may not realize it, but fluorescent light bulbs generally contain small amounts of mercury. Because mercury vapor can be toxic, breathing the gas released when a fluorescent bulb breaks can be dangerous. So used fluorescent bulbs should be disposed of properly and not just placed in landfills with ordinary trash. In view of these facts, we might ask why the mercury is there. Like all elements, mercury has a characteristic way of absorbing and emitting light. The particular pattern of wavelengths absorbed and emitted by mercury is called its atomic spectrum. The existence of such spectra has been known for more than a century. What is most peculiar is that excited atoms of any particular element emit only specific frequencies. Because these frequencies tend to be well separated, the spectra are said to be discrete. This behavior is quite distinct from the light given off by the Sun or by an incandescent bulb, which contains a continuous range of wavelengths. Some surprising observations emerged from investigations of the spectra emitted by various atoms. All elements display discrete spectra, with relatively few wavelengths appearing. The wavelengths that are emitted vary from one element to

The plural of spectrum is spectra.

6.3

Atomic Spectra

169

Gas discharge tube containing hydrogen

Prism

Figure 6.9 ❚ When an electric discharge is applied to a sample of hydrogen gas, H2 molecules dissociate into atoms, and those atoms then emit light. That light passes through a set of slits to create a narrow beam of light, which then passes through a prism. The prism separates the light into its constituent colors and shows that only a few discrete wavelengths are emitted.

the next. Discrete spectra were first observed in the late 19th century and posed an enormous challenge for scientists to interpret. Although data were gathered from many elements, including mercury, the spectrum of hydrogen played the most important role. Hydrogen gas can be excited in a glass tube by passing an electric arc through the tube (Figure 6.9). The energy from the arc breaks the H 2 molecules into individual hydrogen atoms and excites some of those atoms. The excited hydrogen atoms then emit light as they relax toward lower energy states. The light given off appears blue, and when passed through a prism, it reveals four lines visible to the eye. What conclusions can be drawn from the fact that hydrogen (or any other atom) emits only certain wavelengths of light? When an atom emits light, it is releasing energy to the surrounding world. So we should think about this situation in terms of the conservation of energy. When light is emitted, the atom goes from a higher energy state to a lower energy state, and the emitted photon carries away the energy lost by the atom. The fact that only a few wavelengths of light are emitted from a particular atom is direct evidence that the atom can exist in only a few states with very specific energies. To see that this must be true, imagine what would happen if the atom could take on any arbitrary energy. In that case, we would expect to see a continuous spectrum of wavelengths emitted because each atom could lose an arbitrary amount of energy. The notion that an atom could take on only certain allowed energies is extremely counterintuitive because no ordinary macroscopic objects seem to behave similarly. To explain the observed discrete emission spectra, scientists had to develop a model for the behavior of atoms that is fundamentally different from the ways that we describe macroscopic objects.

E XAM PL E PRO BLEM

6 .4

The first four energy levels for the hydrogen atom are shown in the diagram on the next page. The lowest energy state is labeled E1, and the higher states are labeled E2, E3, and E4, respectively. When a hydrogen atom undergoes a transition from E3 to E1, it emits a photon with l = 102.6 nm. Similarly, if the atom undergoes a transition from

170


E3 to E2, it emits a photon with l = 656.3 nm. Find the wavelength of light emitted by an atom making a transition from E2 to E1. E4 E3 λ = 656.3 nm E2 E λ = 102.6 nm

E1

Strategy The energy-level diagram shows us the relationship between the various energy levels involved. The vertical direction in the diagram is energy. So, if we represent each transition by an arrow, as shown, then the length of that arrow will be proportional to the energy lost by the atom in that transition. From the diagram, we can see that the lengths of the arrows for the E3 to E2 transition and for the E2 to E1 transition must add up to the length of the arrow for the E3 to E1 transition. (Another way of thinking about this is that an atom going from the E3 state to the E1 state must lose the same amount of energy whether or not it stops at E2 along the way.) Because we know the wavelengths for two of the transitions, we can find the corresponding photon energies. That will let us find the energy for the third transition, and then finally we can convert that to wavelength.

Solution We first convert the given wavelengths to energies. For the E3 : E1 transition: (6.626 × 10−34 J s)(2.998 × 108 m s−1) hc 109 nm E3:1 = —— = ————————————————— × ———— = 1.936 × 10−18 J l 1m 102.6 nm For the E3 : E2 transition: (6.626 × 10−34 J s)(2.998 × 108 m s−1) hc 109 nm E3:2 = —— = ————————————————— × ———— = 3.027 × 10−19 J l 1m 656.3 nm From the diagram, we can see that E3:1 = E3:2 + E2:1 So, E2:1 = E3:1 − E3:2 = 1.936 × 10−18 J − 3.027 × 10−19 J = 1.633 × 10−18 J Now we just need to convert from energy to wavelength to reach the desired answer. (6.626 × 10−34 J s)(2.998 × 108 ms−1) hc l2:1 = ——— = ————————————————— = 1.216 × 10−7 m E2:1 1.633 nm × 10−18 J Because we know that wavelengths are most often given in nanometers, we might choose to convert that to 121.6 nm.

Analyze Your Answer Our result lies between the two wavelengths we were originally given, which is consistent with the spacing of the energy levels in the diagram. Check Your Understanding A hydrogen atom undergoing a transition from the E4 level to the E3 level emits a photon with a wavelength of 1.873 μm. Find the wavelength of light emitted by an atom making the transition from E4 to E1.

6.3

Atomic Spectra

171

The reasoning used in this example clearly allows us to do some potentially useful calculations based on data from atomic spectra. But it does not yet begin to allow us to explain or understand the very existence of such spectra. In the early part of the 20th century, many models were offered to explain the physics behind atomic spectra. Of these, the model proposed by Niels Bohr had the greatest influence on our subsequent understanding of the atom.

The Bohr Atom Bohr won the Nobel prize for physics in 1922.

For molecules, the concepts of quantized energies and excited states can be extended to vibrational and rotational motions. But we will consider only electronic energy levels.

As the structure of atoms was being deduced, experiments suggested that the nucleus of the atom was positively charged, whereas electrons surrounding that nucleus have a negative charge. You probably learned this model of the atom early in your time as a student, perhaps even in grade school, and so you are likely to accept it without much thought. Yet when first proposed, it was perplexing. In particular, the attraction between oppositely charged particles would seem to indicate that the electrons should collapse into the positively charged nucleus. Bohr addressed this concern as well as other more subtle ones in devising his model for the atom. His most creative contribution was to suggest that the electrons occupy stable orbits from which they cannot deviate without first absorbing or emitting energy in the form of light. The supposition of such orbits had never been made before, but through this model, Bohr was able to explain observations no previous hypothesis could. The influence of the Bohr model is evident in modern society, as the atom is still commonly depicted with electrons orbiting around the nucleus. Now we know that the Bohr model for the atom is not completely correct, but it does provide useful intuition as well as a convenient way to visualize the process of exciting an atom. In any atom, the oppositely charged electrons and protons attract each other. Therefore, the closer to the nucleus an electron is, the stronger the attraction and the lower the energy. For an electron to move from an inner orbit to an outer orbit, as depicted in Figure 6.10, energy must be absorbed. The new distribution or configuration of electrons consists of a grouping that is not at the lowest possible energy and is therefore referred to as an excited state. One of the most powerful driving forces in nature is the tendency to move toward states of lower energy. So an excited state of an atom cannot exist indefinitely. It will return to the lowest energy state, or ground state, by emitting radiation. Both the excitation and emission processes are understandable in terms of the Bohr model of the atom. So this model has some usefulness even though it is not really an accurate representation of atomic structure.

n=1 n=2 n=3 n=4

Figure 6.10 ❚ In the Bohr model, electrons move between allowed orbits when an atom absorbs or emits light. Important elements of this model, including the idea of fixed orbital paths, are not correct. But it played an essential role in the development of our understanding of atomic structure. Note the similarity between this figure and the energy level diagram we used in Example Problem 6.4; both reflect the concept that electrons in atoms are restricted to certain allowed energies.

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This picture of the atom allows us to begin answering some of the questions about light bulbs that we posed in the insight section. Recall that we said that all lamps are based on the principle that an excited substance can emit light as it loses energy. Now we can see that the nature of the excitation largely dictates the type of bulb. In an incandescent bulb, the material excited is a metal wire that is visible to the naked eye. This wire, like any visible material, contains many atoms—perhaps nearly a mole. As electric current passes through the wire, resistance causes the wire’s temperature to rise. Thus the excitation process in the metal wire affects large numbers of atoms by heating them. Emission of energy from the macroscopic collection of excited atoms allows an incandescent bulb to provide us with the light we need. We also noted that fluorescent lamps are usually filled with a rare gas and small amounts of mercury vapor. Remember that the particulate level characteristic of a gas is the relatively large distance between molecules. So, in a fluorescent lamp, the atoms in the bulb are excited individually rather than collectively. Of course, many individual atoms are excited simultaneously, but the behavior of any given atom does not depend on the other atoms in the tube. In a fl uorescent bulb, mercury atoms are excited by an electric arc and then release energy by emitting light. The light emitted by the mercury atoms is in the ultraviolet region of the spectrum, so it cannot be the light that we see when we switch on the light. Instead the phosphor molecules on the glass walls absorb the emitted UV light, and another excitation and emission process takes place. In this second step, the phosphor molecules emit light across a range of frequencies, providing lighting more akin to the nearly continuous light of the Sun, rather than the invisible UV radiation emitted by mercury atoms. The spectrum of wavelengths emitted by the phosphor molecules is not quite the same as that emitted by the hot fi lament in an incandescent bulb, though, and that’s why incandescent and fl uorescent bulbs each have their own characteristic looks.

6.4

The development of improved phosphors allows compact fluorescent bulbs to more closely mimic the light from a standard soft white incandescent bulb.

The Quantum Mechanical Model of the Atom

The observation of atomic spectra and the development of the Bohr model were important advances in the understanding of atoms. But further experiments in the early part of the 20th century revealed still more unexpected properties of atoms, and important limitations of the Bohr model became apparent. So efforts to develop a more robust model of the atom continued. Although these efforts were initially driven by fundamental scientific curiosity, they eventually facilitated many important practical applications. The development of light sources such as the halogen lamp, the light emitting diode (LED), and lasers would not have been possible without improved models of atoms. The picture that emerged to enable these technological advances is the quantum mechanical model of the atom. As we’ll see in Chapter 7, this new model also allows us to describe chemical bonding more accurately, which in turn allows for the design of improved materials. Because the quantum mechanical picture of the atom is both complex and nonintuitive, we begin by pointing out some key comparisons between it and the more familiar Bohr model. In the Bohr model, electrons are viewed as particles traveling along circular orbits of fixed radius. In the quantum model, electrons are viewed as waves rather than particles, and these waves are considered to be spread out, or delocalized, through a region of space called an orbital. The most important similarity between the two models is that in either case the energy of an electron is quantized, meaning it is restricted to certain allowed values. A detailed treatment of the quantum mechanical model is well beyond our grasp here, but in the rest of this section, we will present some of the conceptual and mathematical ideas from which the model arises. As we do

6.4 The Quantum Mechanical Model of the Atom

173

Electron diffraction techniques are now commonly used to probe the structures of solid surfaces.

this, it will help if you keep in mind the simple comparison with the Bohr atom we’ve made here. The essential and peculiar feature of the quantum mechanical model of the atom lies in its description of electrons as waves rather than particles. It is far more intuitive to think of electrons as particles, perhaps resembling tiny marbles, than to envision them as waves. But just as we’ve seen for light, experimental observations led to the idea that electrons can exhibit wave-like behavior. The fi rst evidence of the wave nature of electrons came through diffraction experiments in 1927. Diffraction was already a well-understood phenomenon of waves, so the observation of electron diffraction strongly suggested the need for a wave-based treatment of the electron. A mathematical description of a wave generally begins with a function that describes the periodic behavior. Such a function is referred to as a wave function, for fairly obvious reasons. The simplest such functions are sin x and cos x, which can be used to describe simple oscillatory motion. Although the functions that describe electron waves tend to be more complicated, the idea is the same. The notion of using the mathematics of waves to treat electrons was first put forward by Erwin Schrödinger. The so-called Schrödinger equation is summarized as Hc = Ec

(6.4)

but this form is deceptively simple. That’s because the H term is an operator, which designates a complicated series of mathematical operations to be carried out. E is the energy and c (the Greek letter psi) is the wave function for the electron. If we write out the entire operator H, we find that the Schrödinger equation is actually a secondorder differential equation. (See MathConnections box.) It turns out that in most cases the equation cannot be solved exactly, so the study of quantum mechanics invariably involves approximation techniques. Because of the complexity of the mathematics involved, we will not make any attempt to solve this equation. But we will describe the results of its solution and look at the way these results form the foundation of our modern understanding of atoms. The concept of the wave function, in particular, is central to the quantum mechanical model of the atom. The wave function contains all the information we could ever know about the electron, if we can determine the means by which to extract the information.

MathConnections Written as Hc = Ec, the Schrödinger equation appears deceptively simple. If we write out the operator H for a hydrogen atom, we get the more intimidating looking form below: ≠2c ≠2c ≠2c h2 • —— + —— + —— ¶+ V(x,y,z) c = Ec 2 ——— 8p2m ≠x2 ≠y2 ≠z2 Although this is a complicated equation, we can relate it to things with which you are probably familiar. First, let’s realize that the purpose of this equation is to describe the total energy of an electron as it moves through the three-dimensional space surrounding the atom’s nucleus. Because the nucleus is the center of the atom, it makes sense to define this as the origin of our x, y, z coordinate system. We might expect that there should be terms for the kinetic energy and for the potential energy, and in fact there are. The V(x,y,z) term accounts for the potential energy, which is a function of position. The other term on the left-hand side accounts for the kinetic energy. So the Schrödinger equation expresses the total energy as the sum of the kinetic energy and the potential energy. Next let’s look more closely at the terms that appear. If you are familiar with calculus, you might recognize the “d2/dx2” notation as a derivative. Here it is

174


written with deltas (d) rather than with d’s because x (or y or z) is only one of several variables on which our wave function c depends. The superscript 2’s tell us this is a second derivative. So for d2c/dx2, we would take the derivative of c with respect to x and then take the derivative of the result with respect to x as well. As long as we know how to take derivatives, we could evaluate the kinetic energy term. The potential energy term is also understandable. For an electron in an atom, the potential energy results primarily from the attraction between the positive charge on the nucleus and the negative charge on the electron. So the function V in the equation would represent the way that potential varies as the position of the electron changes. From physics, you may know this as the Coulomb potential, which depends on the distance between charges. You will not be expected to work directly with the Schrödinger equation in a class at this level. But if we think about it step by step like this, you should begin to see the connections between this complicated equation and ideas you work with in your math and physics classes.

Potential Energy and Orbitals The Schrödinger wave equation describes the energy of electrons. As for any object or particle, that total energy includes both kinetic and potential energy contributions. But the potential energy is the most important in describing the structure of atoms. The potential energy of electrons in atoms is associated with the coulombic attraction between the positively charged nucleus and negatively charged electrons as well as the repulsion between like-charged electrons. The periodic nature of waves allows many solutions to the wave equation for any given potential interaction. The Schrödinger equation is often rewritten as Hcn = Encn

We introduced the Coulomb potential in Section 2.3, page 38.

(6.5)

to emphasize the fact that multiple solutions can exist. Here n is an index that labels the different solutions. The wave functions, cn, that provide solutions to this equation for an atom can be written in terms of two components: a radial component, depending only on the distance of the electron from the nucleus, and an angular component that may depend on the direction or orientation of the electron with respect to the nucleus. Just like the simple sin x function, these wave functions may have positive and negative signs in different regions. The most physically meaningful way to interpret a wave function is to look at its square, c 2. This quantity is always positive and tells us the probability of finding an electron at any particular point. This probability interpretation will seem odd to you, but that’s because you are more accustomed to dealing with discrete objects (or particles) rather than with waves. If you think about the standing wave that exists in a vibrating guitar string, you will realize that the wave is not at any one point along the string; the wave is “found” all along the string simultaneously. That’s the same idea we need to apply to electrons in the quantum mechanical model. The indexing system that is represented by the “n” subscript in Equation 6.5 provides important language for discussing atomic structure. The index is typically written to contain both a number and a letter, and these labels identify solutions to the wave equation. Because of their historic correspondence to the position of electrons in the orbits proposed by Bohr, each solution is referred to as an orbital. An orbital is the quantum mechanical equivalent of the ‘location’ of an electron. But because we are treating the electron as a wave, this location is actually a region of space rather than a particular point. It’s quite possible you have heard the names of these orbitals before: 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, etc. What is the origin of these “names” of orbitals, and what information do those names provide?


175

Quantum Numbers Although we don’t need detailed treatments of quantum mechanics to understand light bulbs or most introductory chemical concepts, we will find that the vocabulary established by the theory is important. For example, the names attached to atomic orbitals come from the functions that solve the wave equations. Collectively, they are referred to as quantum numbers. Before examining the specific quantum numbers we need to describe an atomic orbital, let’s look at a simpler problem that can help show where these numbers come from. Suppose we take a guitar string, pin down both ends of it so they can’t move, and then think about what sorts of waves might be produced if we pluck the string so it vibrates (see Figure 6.11). In the simplest case, the length of the string would just be one-half the wavelength of a sine function. In mathematical terms, we could write the following equation for this wave (A is the amplitude of the wave, x is the position along the string, and L is the length of the string): px c(x) = A sin —— L

(6.6)

Other waves are also possible without moving the ends of the string. If a guitarist wants to play what is called the harmonic, a finger is placed very lightly at the midpoint of the string. The string can still vibrate on both sides of the midpoint, so now we can form a sine wave where the length of the string is one complete wavelength, as represented by the following equation. 2px c(x) = A sin ——— L

(6.7)

This wave has a higher frequency, and we would hear it as a higher pitched note. Our guitarist could also produce additional waves of different frequencies by pinning down appropriate positions along the string. The important restriction is that the ends of the string must remain fixed; this means that the length of the string must contain an integral multiple of half-wavelengths. Figure 6.11 ❚ A string fixed at both ends—as in a guitar—is an example of how multiple waves can satisfy a particular set of conditions. In the upper figure, the length of the string is equal to one-half of the wavelength of the wave. Additional cases where the length of the string is equal to one wavelength or to one and one-half wavelengths are also shown. A standing wave can be formed as long as the length of the string is equal to an integral multiple of half wavelengths. But waves that do not satisfy this condition are not possible because they would violate the condition that the ends of the string stay fixed.

String length L = λ/2

String length L = λ

String length L = 3λ/2

String length L = 5λ/4 not possible

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Suppose that we wanted to write a general equation for the waves that can be formed in our string. Working from the cases above, you should be able to see that the necessary form will be npx cn(x) = A sin ——— L

(6.8)

where n must be a positive integer. This n is a quantum number. If we substitute an allowed (positive integer) value for n in the general equation, we will get an equation for a particular wave that can exist in the string. When we solve the Schrödinger equation for an atom, the resulting wave functions are much more complicated than these sine waves. To write the mathematical equations for the wave functions of the atomic orbitals, we need to use three different quantum numbers, which will be described individually below: the principal quantum number (designated n), the secondary quantum number (designated ,), and the magnetic quantum number (designated m,). But just as in our vibrating string problem, these quantum numbers are a natural consequence of the mathematics of the wave equation. Furthermore, when we solve the wave equation for the atomic orbitals, we find that the values which these three quantum numbers are allowed to take on are all interrelated. As we’ll see, this has significant consequences for the kinds of orbitals that exist. The principal quantum number defines the shell in which a particular orbital is found and must be a positive integer (n = 1, 2, 3, 4, 5, . . .). When n = 1, we are describing the first shell; when n = 2, the second shell; and so on. Because there is only one electron in hydrogen, all orbitals in the same shell have the same energy, as predicted by the Bohr model. However, when atoms have more than one electron, the negative charges on the electrons repel one another. This repulsion between electrons causes energy differences between orbitals in a shell. This difficulty was one reason the Bohr model had to be replaced. The secondary quantum number provides a way to index energy differences between orbitals in the same shell of an atom. Like the principal quantum number, it assumes integral values, but it may be zero and is constrained to a maximum value one less than n: , = 0, 1, 2, . . . , n − 1. The secondary quantum number specifies subshells, smaller groups of orbitals within a shell. When n = 1, the only possible value of , is zero, so there is only one subshell. When n = 2, , can be either zero or one, so there are two subshells. Each shell contains as many subshells as its value of n. The secondary quantum number also provides the key to the letter designations of orbitals. Table 6.1 provides the letter designation of the first five values of the , quantum number. The origin of these letters is from the language used to describe the lines seen in early studies of atomic spectra: s was “sharp,” p was “principal,” d was “diffuse,” and f was “fundamental.” Beyond f, the letter designations proceed in alphabetical order. There is no known atom that has electrons in any subshell higher than f when it is in its ground state, but higher orbitals may be occupied in excited states, and additional “superheavy” elements in which these orbitals are populated may yet be discovered. For atoms under normal conditions, the energies of the orbitals are specified completely using only the n and , quantum numbers. In the laboratory, it is possible to place atoms in magnetic fields and observe their spectra. When this is done, some of

Table

❚

6.1

Letter designations for the secondary quantum number ,-value

0

1

2

3

4

Letter Designation

s

p

d

f

g


177

Table

❚

6.2

Relationships among values of the different quantum numbers are illustrated. This table allows us to make another observation about quantum numbers. If we count the total number of orbitals in each shell, it is equal to the square of the principal quantum number, n2.

Value of n

Values for , (letter designation)

Values for m,

Number of Orbitals

1

0 (s)

0

1

2

0 (s) 1 ( p)

0 −1, 0, 1

1 3

3

0 (s) 1 ( p) 2 (d)

0 −1, 0, 1 −2, −1, 0, 1, 2

1 3 5

4

0 (s) 1 ( p) 2 (d) 3(f)

0 −1, 0, 1 −2, −1, 0, 1, 2 −3, −2, −1, 0, 1, 2, 3

1 3 5 7

the emission lines split into three, five, or seven components. This observation points to the need for a third quantum number. This index, called the magnetic quantum number, is labeled m, and is also an integer. It may be either positive or negative, and its absolute value must be less than or equal to the value of , for that orbital. Thus when , = 0, m, is also zero. When , = 1, m, may be −1, 0, or +1, and when , = 2, m, may be −2, −1, 0, +1, or +2. The relationships among the three quantum numbers are summarized in Table 6.2.

E XAM PL E PRO BLEM

6 .5

Write all of the allowed sets of quantum numbers (n, ,, and m,) for a 3p orbital.

Strategy The 3p designation tells us the values of the n and , quantum numbers, and we can use the relationship between , and m, to determine the possible values of m,.

Solution The “3” in “3p” tells us that the orbital must have n = 3, and the “p” tells

us that , = 1. Possible values for m, range from −, to +,; in this case, that gives us , = −1, 0, or +1. So we have three possible sets of quantum numbers for a 3p orbital: n = 3, , = 1, m, = −1; n = 3, , = 1, m, = 0; and n = 3, , = 1, m, = 1.

Check Your Understanding An orbital has quantum numbers of n = 4, , = 2, m, = −1. Which type of orbital (1s, etc.) is this?

Visualizing Orbitals Table 6.3 lists the actual wave functions for the first few orbitals of a hydrogen atom. (Don’t worry—you are not expected to memorize these functions! They are given only to show that they are made up of several standard mathematical functions with which you are already familiar.) Because we expect the atom to be spherical, it is customary to write these functions in spherical polar coordinates, with the nucleus of the atom located at the origin. You are probably familiar with two-dimensional polar

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Table

❚

6.3

Wave functions for the n = 1 and n = 2 shells of a hydrogen atom are listed. There is no need to memorize these functions— they are shown merely to point out that they are made up of ordinary mathematical terms.

Orbital

Wave Function

1s

c1s = p−1/2a−3/2 e−(r/a0) 0

2s

r −(r/2a0) c2s = (4p)−1/2(2a0)−3/2 q2 − —— a re

2px

r −(r/2a0) c2p = (4π)−1/2(2a0)−3/2 —— sinu cosf a e

2py

r −(r/2a0) c2py = (4π)−1/2(2a0)−3/2 —— sinu sinf a e

2pz

r −(r/2a0) c2p = (4π)−1/2(2a0)−3/2 —— cosu a e

The a0 appearing in these equations is a constant called the Bohr radius, with a value of 53 pm.

0

0

x

0

z

0

coordinates, where a point in a plane is specified by a radius (r) and an angle (u). Spherical polar coordinates are the three-dimensional equivalent. To specify a point in three-dimensional space, we need a radius (r) and two angles (u and w), as shown in Figure 6.12. So, the terms that contain r determine the size of the orbital, and those that contain u or w determine the shape. Most chemists usually think of orbitals in terms of pictures rather than mathematical functions and quantum numbers. The pictorial depiction of orbitals requires an arbitrary decision. The wave function provides information about the probability of finding an orbital in a given region of space. If we wish to have a 100% probability of finding an electron, the amount of space needed is quite large and accordingly fails to provide insight about the tendency of that electron to be involved in chemical bonding. Conventionally, the space within which there is a 90% probability of finding an electron is chosen as a boundary for depicting orbitals. Within this convention, a plot of the angular part of the wave function will show us the shape of the corresponding orbital. Pictures of prototype s, p, and d orbitals are shown in Figure 6.13. The s orbital is spherical, whereas the p orbital has two lobes and the d orbital has four lobes. The two lobes of the p orbital are separated by a plane in which there is no probability of finding an electron. This plane is called a node or a nodal plane. The d orbital has two nodes. These nodes give rise to one of the more intriguing philosophical questions of science. If the electron can be on either side of the node, but never in the nodal plane, how does it get from one side to the other? We can suggest an answer to this question in two ways. One is to remember that electrons behave as waves and even classical wave motion such as in the string of a guitar can have nodes.

z

P θ

r y

φ

x

Figure 6.12 ❚ A point P can be specified using three coordinates. Because atoms are spherically symmetric, the spherical polar coordinates r, u, and w, are often used rather than the more familiar x, y, z-coordinate system.

Figure 6.13 ❚ Shapes of (from

2s

2pz

3dyz

left to right) 2s, 2p, and 3d orbitals are shown. Although we sometimes think of figures like these as “pictures,” they are obtained by plotting the corresponding wave functions, as described in the text. The different colored shadings indicate regions in which the wave function has opposite signs.


179

Figure 6.14 ❚ Cross-sectional views of the 1s, 2s, and 3s orbitals are shown. In this type of figure, the density of dots indicates the probability of finding an electron. The greatest electron density will be found in those areas that appear darkest; the white regions within the 2s and 3s orbitals represent nodes.

The actual degree of uncertainty demanded by the uncertainty principle is exceedingly small, so its effects are not detectable for macroscopic objects.

Figure 6.15 ❚ Electron density plots for 3s, 3p, and 3d orbitals show the nodes in each orbital. Although this figure is a sketch, accurate plots can be calculated from wave functions.

n=1 ᐉ=0 ls

The 1s orbital has no nodes.

n=3 ᐉ=0 3s

n=2 ᐉ=0 2s

The 2s orbital has one node.

The 3s orbital has two nodes.

The second and more technical version of the answer lies in a peculiar concept of quantum mechanics called the uncertainty principle. This concept, stated by Werner Heisenberg, expresses the difficulty of observing an electron. Simply put, the uncertainty principle says it is impossible to determine both the position and momentum of an electron simultaneously and with complete accuracy. The uncertainty principle imposes a fundamental limit on observing the microscopic world and forces us to accept the idea that orbitals provide nothing more than probability of finding an electron. The shapes in Figure 6.13 do not tell us anything about the way the probability of finding the electron varies with distance from the nucleus. For that, we need to consider the radial part of the wave function. Depicting this while also considering the angular portion can be somewhat challenging. Figure 6.14 shows cross-sectional “slices” through three different s orbitals corresponding to the principal quantum numbers n = 1, 2, and 3. Note that there are nodes in the radial portion of the wave function just as there are in many of the angular functions (p, d, and f orbitals). There is one less node than the value of the quantum number. Thus there is one node in the 2s orbital, two in the 3s orbital, and so on. Some p and d orbitals, such as 3p, 4p, and 4d, also have radial nodes. Figure 6.15 shows pictures of orbitals in the n = 3 shell. y The 3s orbital has two spherical nodes.

x

y y

x

The 3p orbital has one planar node (right) and a second hyperbolic node (left).

180


x

The 3d orbital has two planar nodes.

Tungsten that evaporates from the filament encounters a halogen atom.

I I

W I Tungsten filament W

I

Figure 6.16 ❚ Traces of halogen gas (here iodine) allow evaporated tungsten atoms to return to the surface of a bulb’s filament, effectively recycling the metal atoms. This allows the filament to operate at a higher temperature without compromising bulb lifetime and explains why halogen lamps are brighter than regular incandescent bulbs. (The chemistry involved is actually much more complicated than shown here.)

The tungsten halide migrates back to the filament and dissociates, replenishing the tungsten and the halogen.

How did something as seemingly abstract as the quantum mechanical model of the atom help the development of modern materials and light sources? The concept of orbital shapes and energies ultimately provided ways to predict chemical reactivities and optimize conditions in which reactions occur. Halogen lamps are predicated on the chemical reaction of halogen molecules with the tungsten wire of the lamp; the role of the halogens is depicted schematically in Figure 6.16. In a normal incandescent bulb, the choice of filament temperature involves a trade-off between the brightness and lifetime of the lamp. Higher temperatures give brighter light but lead to shorter lifetimes because metal atoms evaporate faster from the filament. If halogen gas is present in the bulb, it can react with the tungsten as it vaporizes from the high temperature filament. The resulting molecules redeposit on the wire, recycling the tungsten. These cyclic reactions allow higher operating temperatures and a brighter light without compromising on the lifetime of the bulb. We will not detail the orbital arguments used to understand this process, but the quantum mechanical model of halogens provides important clues. We must, therefore, know more about the way electrons occupy orbitals in atoms such as halogens.

6.5


According to the quantum mechanical model of the structure of atoms, each electron in an atom can be described as occupying a particular orbital. To try to understand the chemical behavior of atoms of different elements, then, we might try to understand how electrons are distributed in orbitals. Let’s begin by asking a very basic question: How many electrons can occupy an orbital? The answer requires the introduction of one more quantum number, the spin quantum number (designated ms). When placed in a strong magnetic field, like those used in magnetic resonance imaging (MRI), electrons behave like tiny magnets. When electric charges move in loops they create magnetic fields. So scientists postulated that electrons (which are negatively charged) must be spinning. This postulate gave rise to the idea and the name of the spin quantum number. Because our modern model describes the electron as a delocalized wave, the spinning particle picture is no longer considered literally correct. But the magnetic properties of electrons are consistent with the notion that they are spinning in one of two directions, and we often refer to electrons as

6.5

In MRI, the “magnets” are protons in hydrogen atoms rather than electrons.


181

The specific values of + or −1/2 are related to the magnitude of the magnetic moment of the electron.

being either “spin up” or “spin down.” We use the quantum number ms to distinguish between these two possible spins, and so it can assume only two possible values, +1/2 or −1/2. Using this quantum number, now, we can specify any electron in an atom by using a set of four quantum numbers. We are constrained, however, in the way we can specify the quantum numbers. We’ve already seen that the values of n, ,, and m, must conform to certain rules that determine which orbitals exist. With the introduction of the ms quantum number, we also need an additional restriction. The Pauli exclusion principle states that no two electrons in an atom may have the same set of four quantum numbers. (Loosely speaking, this is the wave equivalent of saying that two particles cannot occupy the same space.) If two electrons have the same values for n, ,, and m,, then they must have different spin quantum numbers (ms). The major consequence of this principle is that no more than two electrons can occupy any orbital. If two electrons occupy the same orbital, they must be spin paired, one with spin up and the other with spin down. Now that we know that orbitals can hold two electrons, how do we know which orbitals in a given atom are occupied?

Orbital Energies and Electron Configurations We have already noted that in the ground state of an atom, electrons will occupy the lowest energy orbitals available. But how do we know which orbitals are lower in energy, and what causes them to be lower? This question is best considered in two parts. First, realize that a negatively charged electron is attracted to a positively charged nucleus. Electrons in smaller orbitals are held more tightly to the nucleus, so they have lower energies. We’ve already seen from Figure 6.14 that the size of orbitals tends to increase as the value of n increases. Putting these two ideas together, we conclude that the energy of atomic orbitals will increase as the value of the n quantum number increases. For the hydrogen atom, this is the entire story. But for multielectron atoms, we must think about this interaction further. The first difference is the size of the nuclear charge; it will be larger for other elements than for hydrogen. Larger nuclear charges exert stronger attractive forces on electrons, so the size of the orbitals will tend to decrease for higher atomic numbers. More importantly, the electrons in the orbitals interact with other electrons and not only with the positively charged nucleus. For electrons in larger orbitals, therefore, the charge they “feel” is some combination of the actual nuclear charge and the offsetting charge of electrons located in smaller orbitals closer to the nucleus. This masking of the nuclear charge by other electrons is referred to as shielding. Instead of interacting with the full nuclear charge, electrons in multielectron atoms interact with an effective nuclear charge that is smaller than the full charge because of shielding effects. Referring to Figure 6.17, we can see how the concept of effective nuclear charge allows us to understand the energy ordering of orbitals. The graphs show the probability of finding an electron as a function of the distance from the nucleus for each subshell up to n = 3. It is obvious that there are substantial size differences between the shells, such that 1s is notably smaller than 2s and 2p, which in turn are smaller than 3s, 3p, and 3d. This is totally consistent with the argument we just made that orbital energy should increase as n increases. When we consider subshells, however, the distinctions are subtler. Compare the 2s and 2p wave functions. Although the overall sizes of these orbitals are similar, the 2s wave function shows a small local maximum close to the nucleus. No such feature appears in the 2p wave function. This means that a 2s electron has a greater probability of being found very close to the nucleus, and an electron closer to the nucleus will experience a larger effective nuclear charge. As a result, the 2s orbital is lower in energy than the 2p orbitals. Similarly, the 3s orbital, on average, interacts with a larger charge because of the probability it will be close to the nucleus. Even though the first two lobes of the radial wave function are small,

182


1s

n = 1, = 0

3s

n = 3, = 0 r

r

3p

2s

n = 2, = 0 n = 3, = 1 r

r

2p

3d

n = 2, = 1 n = 3, = 2 r

r

Figure 6.17 ❚ Graphs of the radial probability function are shown for hydrogen orbitals with

n = 1, 2, and 3. The increase in orbital size with increasing n is obvious. More subtle variations for orbitals with the same n value are also important in understanding the sequence of orbital energies.

they provide the opportunity for electrons in the 3s orbital to “penetrate” through the shielding of other electrons. This makes the energy of the 3s orbital lower than that of the 3p or 3d orbitals. Similar penetration in the 3p orbitals lowers their energy in comparison to the nonpenetrating 3d orbitals. As we reach the fourth shell of orbitals, the comparison of orbital size to penetrating probability becomes even more difficult. We will not attempt to carry the conceptual reasoning any further, but detailed calculations can be carried out to determine ordering. The accepted order of orbital energies that arises from these calculations is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p. There are no elements currently known whose ground state has electrons that occupy orbitals higher in energy than 7p. You may have learned some device to remember this sequence of orbitals. We will reinforce this ordering in the next section by relating it to the structure of the periodic table. Now that we’ve established the energy sequence of the various subshells, we are ready to decide which orbitals should be occupied in each element. Now we will consider rules for determining these electron configurations.

Hund’s Rule and the Aufbau Principle The process of deciding which orbitals hold electrons in an atom requires that we “fill” orbitals starting with the lowest energy and proceeding to higher energy. In some ways, this concept resembles constructing a building from the ground up, and it is referred to as the aufbau principle. This process is uncomplicated for the first few atoms.

6.5

Aufbau comes from the German word for “build up.”


183

1s

2s

2p

Figure 6.18 ❚ Electron configurations are sometimes depicted using diagrams of the type shown here for carbon. Such diagrams emphasize the presence of unpaired electrons and may be helpful as you become familiar with electron configurations.

The hydrogen atom has one electron, and it occupies the 1s orbital. Helium has two electrons. Both can occupy the 1s orbital, as long as they are spin paired. Lithium has three electrons. Only two can occupy the 1s orbital, and so one electron must go into the next lowest energy orbital, the 2s. A shorthand way to designate this orbital occupation is to write the orbital name followed by a superscript indicating the number of electrons in the orbital. Thus lithium has an electron configuration of 1s22s1. Beryllium is the next element and its atoms contain four electrons. The 2s orbital can also hold two electrons, so the fourth electron is placed there to give a configuration of 1s22s2. With five electrons, boron will also fully populate both the 1s and 2s orbitals, but we must account for an additional electron. The fifth electron goes into the next available orbital, 2p, to give 1s22s22p1. Upon reaching carbon we are faced with a choice. There are three different 2p orbitals, all with equal energy. So two electrons could pair and occupy a single 2p orbital, or they could occupy two different 2p orbitals with one electron in each. The correct choice is motivated by the fact that all electrons are negatively charged particles. Two negatively charged electrons should repel one another, and so the most energetically favorable configuration would have them occupy separate orbitals where they can be farther apart from each other. The concept is summarized in Hund’s rule, which states that within a subshell, electrons occupy orbitals individually whenever possible. To depict this type of distinction in electron configuration requires a device more detailed than the representation we used for the first five elements. One possible choice is to use diagrams with boxes representing orbitals, as shown in Figure 6.18. Another option is to identify individual orbitals with subscripts, such as px, py, and pz. Once we are familiar with electron configurations, though, the simple notation of 1s22s22p2 for carbon will suffice. Knowing Hund’s rule will allow whoever reads this configuration to envision the correct orbital occupation immediately provided by more detailed representations. Practice with this concept is provided in Example Problem 6.6.

E XAM PL E PRO BLEM

6 .6

Sulfur may be an important element in emerging lighting technology. What is the electron configuration of a sulfur atom?

Strategy We use the energy ordering and number of orbitals for each subshell to arrive at the overall electron configuration. Then using Hund’s rule, we can predict the detailed electron configuration for an atom of sulfur.

Solution Sulfur has 16 electrons, so we can place two in 1s, two in 2s, six in 2p, two

in 3s, and four in 3p to get 16. The electron configuration is 1s22s22p63s23p4. For the detailed configuration, we must consider how to distribute the four electrons in the 3p subshell. We imagine placing the first three electrons each into separate 3p orbitals and then filling one of those orbitals by adding the fourth electron. The resulting configuration is shown below. 1s

2s

2p

3s

3p

Check Your Understanding What is the electron configuration of a silicon atom?

184


An observation about fluorescent lights will lead us to a way to abbreviate the writing of electron configurations. Most fluorescent bulbs are filled with argon gas, but some use krypton gas. The fact that these gases are chemically inert is important for their role in the operation of the lamp. What are the electron configurations of these two elements? Argon is 1s22s22p63s23p6 and krypton is 1s22s22p63s23p64s23d104p6. These two elements are chemically similar (both are essentially inert) precisely because they have similar electron configurations. In each case, the last electron fills the p subshell. This correspondence between chemical inertness and a filled p subshell provides the motivation for a shorthand notation for electron configurations. This shorthand notation accomplishes two things. First, it makes writing electron configurations less unwieldy. More importantly, it introduces a very important concept that relates electronic structure to chemical bonding. The electrons in the outermost occupied subshells give rise to the chemical reactivity of an element. Electrons in filled lower shells will have little impact on an element’s chemical properties. The inner electrons, which lie closer to the nucleus, are referred to as core electrons. They can be represented by the atomic symbol of the rare gas element that has the corresponding electron configuration. Thus, although the complete electronic configuration for potassium is 1s22s22p63s23p64s1, the outermost 4s1 electron determines the chemical properties of potassium. We can emphasize this fact by writing the configuration as [Ar]4s1. Similarly, tin (Sn) can be written as [Kr]5s24d105p2. The electrons that are explicitly noted in this notation occupy orbitals that are further from the nucleus. Although it is appropriately descriptive to designate them as outer electrons, they are usually referred to as valence electrons. When we take up the study of chemical bonding in Chapter 7, we will see how important it is to know the valence electron configuration of elements.


We will discuss the connections between electron configuration and chemical bonding in detail in the next chapter.

6.7

Rewrite the electron configuration for sulfur (Example Problem 6.6) using the shorthand notation introduced above.

Strategy We found the electron configuration in Example Problem 6.6. Now we simply need to identify the core electrons and replace them with the symbol for the corresponding noble gas.

Solution The full electron configuration of sulfur is 1s22s22p63s23p4. The valence

electrons occupy orbitals with n = 3, and the core electrons fill the n = 1 and n = 2 shells. Neon is the noble gas element whose electron configuration is 1s22s22p6. So we can abbreviate the configuration for sulfur as [Ne]3s23p4.

Check Your Understanding Use the shorthand notation to write the electron configuration for gallium.

6.6


In talking about lamp technology, we have mentioned two groups of elements, rare gases and halogens. Electron configurations give insight into the roles of these elements in light bulbs. When we first encountered these groups, however, it was in conjunction with the periodic table (in Section 2.5). The relationship between the electron configuration predicted by the quantum mechanical model of the atom and the periodic table was

6.6


185

n 1 1 1s

18 13 14 15 16 17 1s 2p

2

2

2s

3

3s

3

4

4s

3d

5

5s

4d

6

6s

5d

4f

7

7s

6d

5f

Main group elements s block

4

Lanthanides and actinides f block

5

6

7

8

9

10 11 12

3p

3d

4p

4d

5p

5d

6p

6d

7p

Transition elements d block

Main group elements p block

Figure 6.19 ❚ The shape of the periodic table can be broken down into blocks according to the type of orbital occupied by the highest energy electron in the ground state of each element. The f-block elements are usually broken out of the table and displayed separately at the bottom. Here we have placed them in the body of the table to emphasize the sequence in which the various orbitals are filled.

vital to the acceptance of quantum mechanics as a theory. We will take advantage of this important correlation between electron configurations and the periodic table to make it relatively easy to remember configurations for the ground state of any element. The outline of the periodic table in Figure 6.19 includes the elements that are usually shown separately at the bottom. The coloring of regions in this figure divides the table into four blocks of elements. The elements in each block are similar in that the electron in the highest energy orbital comes from the same subshell. Thus all the elements in the section at the left edge of the table, shaded in red, have their highest energy electron in an s orbital. These are sometimes referred to as the s block. The green section at the far right of the table is the p block because according to the aufbau principle, the last orbital occupied is a p orbital. The transition metals are d-block elements (shown in purple), and the lanthanides and actinides make up the f block (shown in yellow). Because of this structure, we can use the periodic table to determine electron configurations for most elements. (A few exceptions arise, mainly among the transition metals.) We find the element of interest in the periodic table and write its core electrons using the shorthand notation with the previous rare gas element. Then we determine the valence electrons by noting where the element sits within its own period in the table. Example Problem 6.8 demonstrates this procedure.

E XAM PL E PRO BLEM

6 .8

Use the periodic table to determine the electron configuration of tungsten (W), which is used in the filaments of most incandescent lights.

Strategy Begin by finding tungsten in the periodic table. Use the rare gas element from the previous period to abbreviate the core electrons. Then determine the correct valence electron configuration by counting across the period in which tungsten is located.

Solution The rare gas element preceding tungsten is xenon. When counting the valence electrons, don’t forget the f-block elements. The correct configuration is [Xe] 6s 24f 145d 4.

Check Your Understanding Mercury is used in fluorescent lamps as well as in mercury vapor lamps. Use the periodic table to determine the electron configuration of mercury (Hg).

186


Emphasizing the connection between the periodic table and electron configuration becomes even more important when we consider chemical reactivity. We’ve seen that halogen lamps are a particularly bright form of incandescent lights. All halogens have electron configurations ending with np5, where n is 2 for fluorine, 3 for chlorine, 4 for bromine, and 5 for iodine. The similarity in electron configuration leads to similarity in chemical behavior and explains why different halogens can be used with similar results. The choice of the specific halogen in a bulb can be adjusted to achieve desired design parameters. But all halogens behave in roughly the same way, enabling a combination of increased brightness and extended lifetime.

6.7


When we first introduced the concept of the periodic table, we noted that various observations about the behavior of atoms led to its development. The reactivity of halogens and the inertness of rare gases are examples of the periodic law that influence the making of light bulbs. Using the understanding of orbitals and atomic structure we have just gained, it is possible to explain some of these periodic variations in greater detail. These atomic properties will then be used to help understand reactivity in later chapters.

Atomic Size One important characteristic of atoms is their size. The measurement of atomic size is sometimes complicated. The results of these measurements are summarized in Figure 6.20. There are two ways to explore trends in the periodic table, within a group and within a period. As we go down a group, we observe an increase in atomic Figure 6.20 ❚ The graph shows atomic radius as a function of atomic number. The different colors and data point symbols represent the different rows in the periodic table. The decrease in size from left to right across each period is clearly seen.

300

250

150

100 Increase Increase

Atomic radius (pm)

200

50

Atomic radii

0 0

10

20

30

40 50 60 Atomic number

70

80

90

6.7

100


187

In Chapter 7 we will use similar arguments to consider the sizes of ions. In general, cations are smaller than their corresponding neutral atoms, whereas anions are larger than their corresponding neutral atoms.

size. As we go across a period, we observe a decrease. An easy way to remember this trend is to draw what look like x, y coordinate axes at the bottom left of the periodic table. An increase in either the x or y direction on these axes corresponds to a decrease in atomic size. How do we understand this trend? The size of the atom is determined largely by its valence electrons because they occupy the outermost orbitals. Two factors are important: (1) the shell in which the valence electrons are found and (2) the strength of the interaction between the nucleus and the valence electrons. As we go down a group in the table, the first of these factors is at work. The size of the valence orbitals increases with n, so the size of the atom increases as we go down the period. The trend in moving across a row is less intuitive. We have already discussed the concept of effective nuclear charge in Section 6.5 (page 182). The distance between a valence electron and the nucleus will tend to be greater if there is a small effective nuclear charge. As we go across a period, we add electrons to the same subshell. These electrons do not shield each other very well, and the effective nuclear charge increases because each increment in atomic number adds a positively charged proton to the nucleus. The higher effective nuclear charge produces a stronger attraction between the electrons and the nucleus. So the valence electrons are drawn closer to the nucleus and the size of the atom decreases.

E XAM PL E PRO BLEM

6 .9

Using only the periodic table, rank the following elements in order of increasing size: Fe, K, Rb, S, and Se.

Strategy We want to arrange the elements from smallest to largest. Begin by finding each element in the periodic table. Then use the general trends for atomic size to place them in the requested order.

Solution Atomic radius increases as we go from top to bottom or right to left in the periodic table. Sulfur and selenium are both in Group 16; because sulfur is above selenium, it will be smaller. Similarly, the positions of rubidium and potassium in Group 1 tell us that potassium is the smaller of the two. Finally, potassium, iron, and selenium are in the fourth period. Selenium is the rightmost of the three, so it will be the smallest, followed by iron and then potassium. Putting all these facts together gives us the requested order: S < Se < Fe < K < Rb.

Check Your Understanding Using only the periodic table, rank the following elements in order of increasing size: Cr, Cs, F, Si, and Sr.

Ionization Energy Another property of atoms that plays a role in the way they interact in forming chemical bonds is the ease with which they lose electrons. This property, called ionization energy, can be measured by an experiment similar to that we described earlier for the photoelectric effect. Ionization energy is defined as the amount of energy we would have to supply to induce the following reaction: X(g) : X+(g) + e− To be more precise, the energy required for this reaction is the first ionization energy. The energy needed to remove an additional electron from the X+ cation, forming X2+, would be the second ionization energy, and so on. We will use the term “ionization energy” for the first ionization energy.

188


Figure 6.21 ❚ The graph shows 2500

He

Increase

First ionization energy (kJ/mol)

Increase

Ne 2000 F

H 1000

Ar

N

1500

C

Be

O

0

10

Zn S

Si Na Al

5

Br

P Mg

Li

Kr

Cl

B 500

First ionization energy

15

the first ionization energy (in kJ/mol) vs. atomic number for the first 38 elements in the periodic table. The inset in the upper right emphasizes the general periodic trend: Ionization energy increases from left to right and bottom to top in the periodic table.

As

Fe Ni Se Cr Ca Ti Ge Co Sr Cu Mn Sc V Ga K Rb

20 25 Atomic number

30

35

Connections between periodicity and atomic structure can be seen by looking at the way ionization energy varies across periods and within groups of the table. Recall from our discussion of atomic size that as we proceed across a period, the effective nuclear charge increases. Not surprisingly, the ionization energy also generally increases across a period. As we move down a group, the valence electrons occupy larger and larger orbitals. As the valence electrons move farther from the nucleus, they become easier to remove. So ionization energy decreases as we move down a column in the periodic table. As we proceed up and toward the right in the periodic table, the ionization energy increases, as shown in Figure 6.21. Additional clues about atomic structure are revealed by the exceptions to the trends in ionization energy. Looking closely at Figure 6.21, we see that from nitrogen to oxygen, there is a slight decrease in ionization energy rather than the increase expected based on the periodic trend. This observation can be rationalized by considering electron pairing. Nitrogen has a half-filled p subshell, with one electron in each 2p orbital. To accommodate its p4 configuration, oxygen must pair two electrons in one of the 2p orbitals. Removing one electron, therefore, leaves only unpaired electrons in all three 2p orbitals. Such a change leads to notably less electron–electron repulsion, enough so that the ionization energy of oxygen is lower than that of nitrogen. Similar exceptions to trends are seen for sulfur in the third period and selenium in the fourth.


6 .10

Using only the periodic table, rank the following elements in order of increasing ionization energy: Br, F, Ga, K, and Se.

Strategy We want to arrange the elements from the smallest ionization energy to the largest. Begin by finding each element in the periodic table. Then use the general trends for ionization energy to place them in the requested order.

Solution Ionization energy increases as we go up or to the right in the periodic table. From our list of elements, all but fluorine are in the fourth period. So we expect their ionization energies to increase from left to right: K < Ga < Se < Br. Finally, fluorine lies above bromine in Group 17, so it will have the largest ionization energy. This gives us the requested order: K < Ga < Se < Br < F.

6.7


189

Check Your Understanding Using only the periodic table, rank the following elements in order of increasing ionization energy: He, Mg, N, Rb, and Si.

We can also gain insight into atomic structure by looking at successive ionization energies for the same element, as summarized in Table 6.4. For sodium, the first ionization energy is among the lowest, but the second is much larger than that for any neutral atom. For magnesium, the situation is quite different. The second ionization is still larger than the first because the effective nuclear charge experienced by a valence electron is larger in the ion than in the neutral atom. But the more dramatic change in ionization energy lies between the second and third ionizations. The point at which this large jump occurs for each element is highlighted throughout the table. In each case, the largest increase in ionization energy occurs when an electron is removed from a completely filled p subshell. From this observation, we can see the unusual stability of that filled p subshell, and we can begin to appreciate why noble gas elements are so unreactive. Filled subshells of electrons are difficult to break up, and as we’ll see, this fact will help us to understand chemical bonding.

Table

❚

6.4

The first four ionization energies (all in kJ/mol) for the elements of the first three periods. Those ionizations with values shown in shaded cells with bold print involve removing the last electron from a particular shell. Further ionization requires removing an electron from a more stable filled shell, and this leads to a very large increase in ionization energy.

190

Z

Element

IE1

IE2

IE3

IE4

1

H

1,312

—

—

—

2

He

2,372

5,250

—

—

3

Li

520.2

7,298

11,815

—

4

Be

899.4

1,757

14,848

21,007

5

B

800.6

2,427

3,660

25,026

6

C

1,086

2,353

4,620

6,223

7

N

1,402

2,856

4,578

7,475

8

O

1,314

3,388

5,300

7,469

9

F

1,681

3,374

6,050

8,408

10

Ne

2,081

3,952

6,122

9,370

11

Na

495.6

4,562

6,912

9,544

12

Mg

737.7

1,451

7,733

10,540

13

Al

577.6

1,817

2,745

11,578

14

Si

786.4

1,577

3,232

4,356

15

P

1,012

1,908

2,912

4,957

16

S

999.6

2,251

3,357

4,564

17

Cl

1,251

2,297

3,822

5,158

18

Ar

1,520

2,666

3,931

5,771


Electron Affinity Ionization energy is defined as the energy required to remove an electron from an atom, forming a positively charged cation. But we also know that some atoms will routinely pick up an electron to form a negatively charged anion. X(g) + e− : X−(g) We can define a property called electron affinity as the amount of energy we would have to supply to induce this reaction. Removing an electron from a neutral atom always requires an input of energy because the atom will not simply eject electrons on its own. So ionization energies are always positive. But electron affinity values can be either positive or negative. If the resulting anion is stable, then energy will actually be released as the electron is added. But our definition of electron affinity says it is the amount of energy that must be supplied to form the anion. If energy is being released, we might say that we must “supply a negative amount of energy,” so the electron affinity value will be negative. If the electron affinity is positive, that means that the anion formed is not stable; in this case, we would need to supply energy to force the atom to accept the electron. Most electron affinity values are negative. This definition of electron affinity may seem a little twisted because it is at odds with the normal meaning of the word “affinity.” Under our definition, if the atom “wants” the electron, its electron affinity value is negative. But we define the signs in this way because it is consistent with the conventions usually followed when discussing energy in chemistry, physics, and engineering. Energy added to an atom is positive, and energy released by an atom is negative. Electron affinity values display periodic trends based on electron configuration, much like the other properties we’ve been considering. The most immediate conclusion from Figure 6.22 might be that electron affinity values vary more erratically than ionization energies. Still, it should be evident from the figure that as we move from bottom to top or from left to right in the periodic table, the electron affinity tends to become more negative. Thus in terms of the x, y scheme we have been using for remembering general trends, electron affinity becomes more negative with increasing x or y, much like ionization energy. It is customary to think of the trend in electron affinity in this way, even though the value is actually becoming algebraically smaller as it gets more negative. That’s because the more negative the electron affinity becomes, the more stable the anion will be. Closer examination of the electron affinity values reveals that there are exceptions in the trends associated with electron pairing, just as there were for ionization energy. Thus, nitrogen does not have a negative electron affinity because adding an electron would force the pairing of electrons in a p orbital, a process that increases the electron–electron repulsion substantially.

Increase (more negative)

He Be

0

B

Al Na

H Li

⫺100

Ar Ca

Ne Mg

N

K

Electron affinity

Si

C O

⫺200

P

Increase

Electron affinity (kJ/mol)

100

S

⫺300 F ⫺400

0

5

10 15 Atomic number

The study of energy changes in chemical reactions is called thermodynamics and is presented in Chapters 9 and 10.

Figure 6.22 ❚ The graph shows the electron affinity (in kJ/mol) vs. atomic number for the first 20 elements of the periodic table. The inset at the upper right emphasizes the general periodic trend: Electron affinity “increases”—meaning that formation of the anion becomes more favorable—from left to right and bottom to top in the periodic table.

Cl 20

25

6.7


191

There are other periodic trends in the properties of the elements that we will not treat explicitly here. Some of those properties help dictate the use of materials for technology, including light bulbs. The choice of tungsten for the filament in incandescent bulbs is partly an economic issue, but it also depends on properties such as ductility and melting point. Because these properties are related to the bulk metal rather than to isolated atoms, they are more complicated than those discussed thus far. But the electron configuration of tungsten does provide some insight. Melting a metal requires breaking metallic bonds. Having more unpaired electrons present in an atom tends to increase the bonding strength of a metal in the solid state. Stronger bonding in the solid metal means that the melting point will be higher. So, the four unpaired d electrons in tungsten help produce a solid that melts at high temperatures, allowing the filament to operate at high enough temperatures to produce a bright light output.

INSIGHT INTO

6.8 Although Thomas Edison is often regarded as the inventor of the electric light bulb, the first patent for an incandescent lamp was granted to Frederick de Moleyns of London in 1841.

Some LEDs emit two or more wavelengths simultaneously, but usually they are so closely spaced that the human eye cannot separate them.

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Modern Light Sources: LEDs and Lasers

The history of the incandescent light bulb dates back to the middle of the 19th century. Over the years, numerous developments have improved the performance and efficiency of incandescent lamps. Advances in our understanding of chemistry have enabled several of these improvements, such as the use of halogen gases. Similarly, developments in chemistry have helped identify new phosphor molecules that drove improvements in fluorescent lamps. Today’s technologies rely heavily on more modern light sources, including light-emitting diodes (LEDs) and lasers. The very existence of these 20th century creations was made possible by improvements in our understanding of atomic structure and chemical bonding. How does the light produced by an LED or a laser differ from that given off by an incandescent or fluorescent lamp? And what particular attributes of these devices and the light they emit make them so well suited for particular tasks? You can almost certainly find light-emitting diodes in many of the appliances you use every day. The power indicators on everything from cell phones to TVs and the displays on many digital alarm clocks and microwave ovens are just a few examples of common uses for LEDs. As you think about some of the applications in which they are used, you may realize one of the characteristics of LEDs: each LED emits light of a particular color. Although you can buy colored incandescent bulbs, they rely on some sort of filter coated onto the surface of the bulb. But the filament still emits white light, so the colored filter greatly reduces the efficiency of the bulb. In an LED, the light emitted is usually monochromatic, which means that it is restricted to a single wavelength or color. So LEDs provide an easy way to produce colored light, and that’s one reason that they are so widely used as indicators on electronic devices. Another important contrast between an LED and a small incandescent bulb is in durability: light bulbs break or burn out, but LEDs last almost indefi nitely under normal conditions. We’ve seen that bulbs burn out because their filaments gradually break down at the high operating temperatures needed to produce light. An LED, though, does not require a high operating temperature. It functions as long as a small electric current flows through it. This gives the LED a long useful lifetime and also reduces the amount of power consumed and heat given off, both of which can be important design considerations in modern electronics. Because they lose very little energy as heat, LEDs can operate at very high efficiency, converting nearly all of the input energy into light. This is especially important for small portable devices, such as cell phones or MP3 players, where battery lifetime is a major consideration.


Lens

Leads

Photo: © Cengage Learning/ Charles D. Winters

Semiconductor

Figure 6.23 ❚ The standard design of an LED is extremely simple. Metallic leads allow a current to pass through a small piece of a semiconductor material, which then emits light. A plastic enclosure serves as both a lens and a case for the device. The color of light emitted is determined by the chemical composition of the semiconductor. Because LEDs are superior to incandescent lights in efficiency and durability, they are now routinely used in traffic lights, like the one shown here.

LEDs are an example of what are called solid-state devices, in which the primary component is simply a piece of some solid material (Figure 6.23). The important functional properties of such devices are determined primarily by the composition of this solid material—and that’s where chemistry comes in. The color of the light emitted by an LED depends on the details of the chemical bonding between the atoms in the solid. As you might guess, that bonding depends on the identity of the atoms involved. So adjusting the composition of the solid can change the color of light emitted. We’ll look at some of the details behind this in Section 8.3. But the practical implication is that we can produce an LED that emits almost any color of light. Lasers share some traits with LEDs: they produce monochromatic light, and the wavelength of that light depends on the specific properties of the laser. But lasers also possess other important properties. Perhaps most obviously, the light emitted by a laser is highly directional and emerges as a well-defined beam. You’ve almost certainly seen laser pointers or other lasers, so you may know that the beam is generally not even visible from the side because all of the light is traveling in the same direction. Laser light is also coherent, which means that all of its waves are perfectly in phase and go through maxima and minima together. Many varieties of lasers are known, ranging from small devices used in DVD players, laser printers, and supermarket price scanners to very high-powered systems used in industrial machining. The invention of the laser is a direct result of the emergence of the quantum mechanical model of the atom. Laser light is emitted as atoms or molecules go from a higher energy level down to a lower energy level, just as we saw in atomic spectra. So every laser’s design is ultimately based on a detailed understanding of the energy levels available to the atoms or molecules in what is called the laser medium. The operation of a laser generally relies on a phenomenon called stimulated emission, which becomes important when the population of one energy level becomes greater than the population of a lower energy level. Many of the most common lasers, like those in your DVD players, are solid-state devices in which this population inversion is established for the electrons in a solid. In other lasers, atoms or molecule in a gas or in solution emit the light. The primary engineering challenges in designing a laser are to identify a medium that can emit the desired wavelength of light and then to devise a way to establish and maintain the population inversion. So an understanding of the energy levels involved is clearly essential. As new lighting technologies emerge in the 21st century, the connections between chemistry and engineering continue to play a major role. Recent improvements in liquid crystal displays used in flat panel computer monitors, for example, are due in part to the development of new chemical compounds whose light absorbing properties are tailored for their tasks. Research into organic light emitting diodes (OLEDs) is moving forward rapidly and holds the promise of an entirely new class of paper-thin, flexible color displays. Advances in all of these areas require the close cooperation of research chemists and design engineers.

For many years blue LEDs were not available because their higher photon energy made them more challenging to produce.

6.8 Modern Light Sources: LEDs and Lasers

193

Courtesy of the U.S. Department of Energy’s Ames Laboratory

The bright blue light in this photo is given off by an organic light emitting diode (OLED). These emerging devices hold great promise for use in small, flexible displays, and versions emitting white light may also find use as light sources.

FOCU S O N

PRO BLEM S O LVI NG

Question (a) Based on your understanding of atoms and their spectra, speculate on why streetlights that use sodium lamps seem yellow/orange in color. (b) What device or instrument could be used to make a measurement that would test your hypothesis from (a)?

Strategy This question involves conceptual problem solving. We can approach questions like this by asking what atomic or molecular level information must be inferred to answer it. Then we need to think about what sort of measurement or experiment could provide the atomic or molecular information needed to confirm or refute our idea.

Answer (a) The key words in this question are “sodium lamp.” This description implies that there is something important about the lamp that is related to the element sodium. We saw in this chapter that atoms have distinctive spectra, so it is reasonable to surmise that in the case of sodium atoms enough colors are missing that the remaining light appears yellow/orange. (b) We could test this hypothesis by using a spectrometer or prism to view the light from the lamp. If our assertion is correct, the spectrum should show lines in the yellow-orange region but not in other parts of the visible spectrum.

S U MMA RY To develop an understanding of matter at the microscopic level, first we must have a sound model for the structure of atoms. Much of our current understanding of atoms grew out of studies of the interaction of atoms with electromagnetic radiation, or light. Many properties of light can be explained by a wave model in which the frequency or wavelength defines the color of the light and the amplitude defines its intensity. But we can only understand some aspects of the interaction of light with atoms by using a particulate model in which we think of light as packets of energy, called

194


photons. We thus encounter the concept of wave-particle duality for the first time when we consider the nature of light. The most direct connection between atoms and light is the observation of discrete atomic spectra. In this context, discrete means that when atoms absorb or emit light they do so at only a few frequencies rather than across the entire spectrum. This observation leads to the hypothesis that the electrons in an atom can occupy only a limited number of energy levels: the energies of electrons in atoms are said to be quantized. Over time, various

models have been introduced and improved to account for these quantized energies. Our best current model is the quantum mechanical theory of the atom. In quantum theory, electrons are treated as waves that can be described by mathematical wave functions. (This is another instance of wave-particle duality.) The detailed mathematics of the quantum mechanical model is responsible for much of the vocabulary we use to describe atomic structure. Thus, electrons occupy orbitals that are characterized by quantum numbers. For the purpose of understanding chemistry, there are two important things to remember about this model. First, orbitals have sizes and shapes that can be correlated with the labeling scheme derived from quantum numbers. Second, for atoms under most ordinary conditions, we can assign the electrons to the lowest energy orbitals available.

These two ideas lead to the concept of electron configurations for atoms—a concept that ultimately will help us gain an understanding of chemical bonding. At this point, we have already seen that the quantum model and electron configuration can explain many facets of the structure of the periodic table. The familiar shape of the periodic table arises from the fact that the s-block elements in Groups 1 and 2 have electron configurations where the highest energy electron is in an s orbital. Similarly, Groups 13–18 comprise the p-block elements because the highest energy electron is in a p orbital. In addition to this overall structure of the periodic table, trends in atomic properties such as size or ionization energy can also be explained by considering the quantum mechanical model of the atom and the electron configurations that arise from it.

K EY TE RM S amplitude (6.2)

Hund’s rule (6.5)

Schrödinger equation (6.4)

atomic spectrum (6.3)

ionization energy (6.7)

secondary quantum number (𝓵) (6.4)

aufbau principle (6.5)

laser (6.8)

shell (6.4)

binding energy (6.2)

light-emitting diode (LED) (6.8)

shielding (6.5)

Bohr model (6.3)

magnetic quantum number (m,) (6.4)

spin paired (6.5)

coherent light (6.8)

monochromatic (6.8)

spin quantum number (ms) (6.5)

core electrons (6.5)

node (6.4)

subshell (6.4)

diffraction (6.4)

operator (6.4)

uncertainty principle (6.4)

effective nuclear charge (6.5)

orbital (6.4)

valence electrons (6.5)

electromagnetic spectrum (6.2)

Pauli exclusion principle (6.5)

visible light (6.2)

electron affinity (6.7)

photoelectric effect (6.2)

wave function (6.4)

electron configuration (6.5)

photon (6.2)

wavelength (6.2)

excited state (6.3)

principal quantum number (n) (6.4)

wave-particle duality (6.2)

frequency (6.2)

quantum number (6.4)

ground state (6.3)

refraction (6.2)



INSIGHT INTO Incandescent and Fluorescent Lights 6.1 How is the color of molten iron (at a blacksmith’s shop, for example) related to the color of a light bulb?

The Electromagnetic Spectrum 6.7 Explain why light is referred to as electromagnetic radiation. 6.8 Which of the waves depicted here has the highest frequency? Explain your answer.

6.2 What is the first step in the process by which an incandescent bulb provides light? 6.3 What material is most commonly used for the filaments in an incandescent bulb? Why is this particular material chosen? 6.4 How is electricity used differently in an incandescent light bulb and in a fluorescent light bulb?

(a)

(b)

6.5 Describe what is meant by the term electric arc. 6.6 What is the role of the electric arc in a fluorescent light bulb?

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(c)


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6.9 Arrange the following regions of the electromagnetic spectrum in order of increasing frequency: IR, UV, radio wave, visible. 6.10

Calculate the wavelengths, in meters, of radiation of the following frequencies: (a) 5.00 × 1015 s–1, (b) 2.11 × 1014 s–1, (c) 5.44 × 1012 s−1 ■

6.11 Decorative lights, such as those found on Christmas trees, often achieve their colors by painting the glass of the bulb. If a string of lights includes bulbs with wavelengths of 480, 530, 580, and 700 nm, what are the frequencies of the lights? Use Figure 6.6 to determine which colors are in the set.

Atomic Spectra 6.25 What is the difference between continuous and discrete spectra? 6.26 What was novel about Bohr’s model of the atom? 6.27 Describe how the Bohr model of the atom accounts for the spectrum of the hydrogen atom. 6.28 According to the Bohr model of the atom, what happens when an atom absorbs energy? 6.29 Define the term ground state. ■

6.31

■ A mercury atom emits light at many wavelengths, two of which are at 435.8 and 546.1 nm. Both of these transitions are to the same final state. (a) What is the energy difference between the two states for each transition? (b) If a transition between the two higher energy states could be observed, what would be the frequency of the light?

6.12 Define the term refraction. 6.13 Define the term photon. 6.14 What is the relationship between the number of photons and the intensity of light? 6.15 Find the energy of a photon with each of the following frequencies. (a) 15.3 THz, (b) 1.7 EHz (see Table 1.2 if needed), (c) 6.22 × 1010 Hz 6.16

■ Place

these types of radiation in order of increasing energy per photon. (a) green light from a mercury lamp, (b) X-rays from a dental X-ray, (c) microwaves in a microwave oven, (d) an FM music station broadcasting at 89.1 MHz

6.17 For photons with the following energies, calculate the wavelength and identify the region of the spectrum they are from: (a) 3.5 × 10−20 J, (b) 8.7 × 10−26 J, (c) 7.1 × 10−17 J, (d) 5.5 × 10−27 J 6.18 For photons with the following energies, calculate the wavelength and identify the region of the spectrum they are from: (a) 6.0 × 10−19 J, (b) 8.7 × 10−22 J, (c) 3.2 × 10−24 J, (d) 1.9 × 10−28 J 6.19

Various optical disk drives rely on lasers operating at different wavelengths, with shorter wavelengths allowing a higher density of data storage. For each of the following drive types, find the energy of a single photon at the specified wavelength. (a) CD, l = 780 nm, (b) DVD, l = 650 nm, (c) Blu-ray® disc, l = 405 nm

■

6.20 The laser in most supermarket barcode scanners operates at a wavelength of 632.8 nm. What is the energy of a single photon emitted by such a laser? What is the energy of one mole of these photons? 6.21

■ Assume that a microwave oven operates at a frequency of 1.00 3 1011 s−1. (a) What is the wavelength of this radiation in meters? (b) What is the energy in joules per photon? (c) What is the energy per mole of photons?

6.22 Fill in the blanks below to complete a description of the photoelectric effect experiment. (You should be able to do this with just one or two words in each blank.) A beam of causing 6.23

6.24

196

strikes

,

to be emitted.

■ When light with a wavelength of 58.5 nm strikes the surface of tin metal, electrons are ejected with a maximum kinetic energy of 2.69 × 10−18 J. What is the binding energy of these electrons to the metal?

The electron binding energy for copper metal is 7.18 × 10−19 J. Find the longest wavelength of light that could eject electrons from copper in a photoelectric effect experiment.

■


A neon atom emits light at many wavelengths, two of which are at 616.4 and 638.3 nm. Both of these transitions are to the same final state. (a) What is the energy difference between the two states for each transition? (b) If a transition between the two higher energy states could be observed, what would be the frequency of the light?

6.30

6.32 Describe how excited states of atoms play a role in light bulb technology. 6.33 Why do fluorescent bulbs have phosphors on the walls of the bulb? The Quantum Mechanical Model of the Atom 6.34 How did the observation of electron diffraction affect the development of the quantum mechanical model of the atom? 6.35 Why do we use a wave function to describe electrons? 6.36 What are the mathematical origins of quantum numbers? 6.37 What are the allowed values for the principal quantum number? For the secondary quantum number? 6.38

■ Which of the following represent valid sets of quantum numbers? For a set that is invalid, explain briefly why it is not correct. (a) n = 3, , = 3, m, = 0, (b) n = 2, , = 1, m, = 0, (c) n = 6, , = 5, m, = −1, (d) n = 4, , = 3, m, = −4

6.39

■A

particular orbital has n = 4 and , = 2. What must this orbital be? (a) 3p, (b) 4p, (c) 5d, or (d) 4d

6.40 Why are there no 2d orbitals? 6.41

What is the maximum number of electrons in an atom that can have the following quantum numbers? (a) n = 2; (b) n = 3 and , = 1; (c) n = 3, , = 1, and m, = 0; (d) n = 3, , = 1, m , = −1, and m s = −1/2.

6.42

■

■

How many orbitals correspond to each of the following designations? (a) 3p, (b) 4p, (c) 4px, (d) 6d, (e) 5d, (f ) 5f, (g) n = 5, (h) 7s

6.43 The p and d orbitals are sometimes referred to as directional. Contrast them with s orbitals to offer an explanation for this terminology. 6.44 Define the term nodal plane (or node). 6.45 Referring to Figure 6.15, draw a 4p orbital, showing all of its nodes.

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The Pauli Exclusion Principle and Electron Configurations 6.46 Does the fact that the fourth quantum number is called the spin quantum number prove that electrons are spinning? What observations gave rise to this name? 6.47 Define the term spin paired. 6.48 On what does the Pauli exclusion principle place a limit? 6.49 Define the term shielding. 6.50 Why does the size of an orbital have an effect on its energy? 6.51 What is effective nuclear charge and how does it influence the size of an orbital? 6.52 Depict two ways to place electrons in the 2p orbitals for a nitrogen atom. Which depiction is correct according to Hund’s rule? 6.53 How does the charge of electrons provide some rationale for Hund’s rule?

6.67

Using only a periodic table as a guide, arrange each of the following series of atoms in order of increasing size. (a) Na, Be, Li, (b) P, N, F, (c) I, O, Sn

■

6.68 Define the term ionization energy. What is the difference between the first ionization energy and the second ionization energy? 6.69 At which ionization for chlorine would you expect the first large jump in ionization energy? Would this be the only large jump in energy if you continued to ionize the chlorine? 6.70

Arrange the following atoms in order of increasing ionization energy: Li, K, C, and N.

■

6.71 How do we explain the fact that the ionization energy of oxygen is less than that of nitrogen? 6.72 Which element would you expect to have the largest second ionization energy, Na, C, or F? Why? 6.73

■ Answer each of the following questions. (a) Of the elements S, Se, and Cl, which has the largest atomic radius? (b) Which has the larger radius, Br or Br –? (c) Which should have the largest difference between the first and second ionization energy: Si, Na, P, or Mg? (d) Which has the largest ionization energy: N, P, or As? (e) Which of the following has the largest radius: O2–, N3–, or F – ?

6.54

■ Write the ground state electron configuration for (a) B, (b) Ba, (c) Be, (d) Bi, (e) Br.

6.55

■ Which of these electron configurations are for atoms in the ground state? In excited states? Which are impossible? (a) 1s22s1, (b) 1s22s22p3, (c) [Ne] 3s23p34s1, (d) [Ne] 3s23p64s33d2, (e) [Ne] 3s23p64 f 4, (f) 1s22s22p43s2

6.74

6.56

■

From the list of atoms and ions given, identify any pairs that have the same electron configurations and write that configuration: Na+, S2−, Ne, Ca2+, Fe2+, Kr, I−

■ Indicate which species in each pair has the more favorable (more negative) electron affinity. Explain your answers. (a) Cl or S, (b) S or P, (c) Br or As

6.75

■ Compare

The Periodic Table and Electron Configurations 6.57 Halogen lamps rely on the predictable chemistry of the halogen elements. Compare the electron configurations of three halogens and use this comparison to justify the similarity of the chemistry of these elements. 6.58 Distinguish between the terms core electrons and valence electrons. 6.59 Describe how valence electron configurations account for some of the similarities in chemical properties among elements in a group. 6.60 Why is there no element to the immediate right of magnesium in the periodic table? 6.61 If another 50 elements were discovered, how would you expect the appearance of the periodic table to change? 6.62 Which blocks of the periodic table comprise the main group elements? 6.63 What is the main difference between a halogen lamp and an ordinary incandescent lamp? 6.64 Look at the table of electron configurations in Appendix C. Which elements have configurations that are exceptions to the aufbau principle? Propose a reason why these elements have these exceptions. Periodic Trends of Atomic Properties 6.65 How does the effective nuclear charge help explain the trends in atomic size across a period of the periodic table? 6.66 Use the electronic configuration of the alkali metals and your knowledge of orbitals and quantum numbers to explain the trend in atomic size of alkali metal atoms.

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the elements Na, B, Al, and C with regard to the following properties. (a) Which has the largest atomic radius? (b) Which has the most negative electron affinity? (c) Place the elements in order of increasing ionization energy.

6.76 Rank the following in order of decreasing ionization energy. Cl, F, Ne+, S, S− INSIGHT INTO Modern Light Sources: LEDs and Lasers 6.77

Several excited states of the neon atom are important in the operation of a helium-neon laser. In these excited states, one electron of the neon atom is promoted from the 2p level to a higher energy orbital.

■

An excited neon atom with a 1s22s22p55s1 electron configuration can emit a photon with a wavelength of 3391 nm as it makes a transition to a lower energy state with a 1s 22s 22p 54p 1 electron confi guration. Other transitions are also possible. If an excited neon atom with a 1s 2 2s 2 2p 5 5s 1 electron configuration makes a transition to a lower energy state with a 1s22s22p53p1 electron confi guration, it emits a photon with a wavelength of 632.8 nm. Find the wavelength of the photon that would be emitted in a transition from the 1s22s22p54p1 electron configuration to the 1s22s22p53p1 electron configuration. (It should help if you start by drawing an energy-level diagram.) 6.78 The compact fluorescent bulb is another fairly new light source that is likely to have a large impact on home lighting. These bulbs combine the energy efficiency of traditional fluorescent lights with the smaller size, “warmer” light, and dimming ability of incandescent lights. Although compact fluorescent bulbs are expensive, manufacturers claim that


197

subsequent savings will more than offset the initial costs. Analyze the relative cost of incandescent versus compact fluorescent lighting, assuming that the information in the table below is accurate. Are industry claims of cost savings justified?

Initial cost Electricity usage Electricity cost Lifetime

Compact Fluorescent

Incandescent

$15.00

$2.00

18 watts

75 watts

$0.04/kwh

$0.04/kwh

10,000 hours

2,500 hours

6.81

6.82

A mercury atom is initially in its lowest possible (or ground state) energy level. The atom absorbs a photon with a wavelength of 185 nm and then emits a photon with a frequency of 6.88 × 1014 Hz. At the end of this series of transitions, the atom will still be in an energy level above the ground state. Draw an energy-level diagram for this process and find the energy of this resulting excited state, assuming that we assign a value of E = 0 to the ground state. (This choice of E = 0 is not the usual convention, but it will simplify the calculations you need to do here.)

■ When a photoelectric effect experiment was carried out using a metal “M” and light at wavelength l1, electrons with a kinetic energy of 1.6 × 10−19 J were emitted. The wavelength was reduced to one-half its original value and the experiment was repeated (still using the same metal target). This time electrons with a kinetic energy of 6.4 × 10−19 J were emitted. Find the electron binding energy for metal M. (The actual wavelengths used were not recorded, but it is still possible to find the binding energy.)

6.83 A metallic sample is known to be barium, cesium, lithium, or silver. The electron binding energies for these metals are listed in the following table.

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4.30 × 10−19

Cesium

3.11 × 10−19

Lithium

3.94 × 10−19

Silver

7.59 × 10−19

Photoelectrons Seen?

532 nm

No

488 nm

Yes

308 nm

Yes

Based on this information, what conclusions can be drawn as to the identity of the metal?

■

■

Barium

Laser Wavelength

Additional Problems The photoelectric effect can be used to measure the value of Planck’s constant. Suppose that a photoelectric effect experi ment was carried out using light with ν = 7.50 × 1014 s−1 and ejected electrons were detected with a kinetic energy of 2.50 × 10−11 J. The experiment was then repeated using light with ν = 1.00 × 1015 s−1 and the same metal target, and electrons were ejected with kinetic energy of 5.00 × 10−11 J. Use these data to find a value for Planck’s constant. HINTS: These data are fictional and will give a result that is quite different from the real value of Planck’s constant. Be sure that you do not use the real value of Planck’s constant in any calculations here. It may help to start by thinking about how you would calculate the metal’s binding energy if you already knew Planck’s constant.

Electron Binding Energy ( J)

One way to identify the element might be through a photoelectric effect experiment. The experiment was performed three times, each time using a different laser as the light source. The results are summarized below. (The kinetic energy of the ejected photoelectrons was not measured.)

6.79 How much energy could be saved each year by replacing all the incandescent bulbs in a typical house with compact fluorescent bulbs? Assume that six 75-watt lamps are lit an average of three hours a night and they are each replaced with 18-watt compact fluorescent bulbs. (See Problem 6.78.)

6.80

Metal

6.84

■ When a helium atom absorbs light at 58.44 nm, an electron is promoted from the 1s orbital to a 2p orbital. Given that the ionization energy of (ground state) helium is 2372 kJ/mol, find the longest wavelength of light that could eject an electron from the excited state helium atom.

6.85

■ Arrange

the members of each of the following sets of cations in order of increasing ionic radii: (a) K+, Ca2+, Ga3+, (b) Ca2+, Be2+, Ba2+, Mg2+, (c) Al3+, Sr2+, Rb+, K+, (d) K+, Ca2+, Rb+

6.86

■ Arrange the following sets of anions in order of increasing ionic radii: (a) Cl−, S2−, P3−, (b) O2−, S2−, Se2−, (c) N3−, S2−, Br−, P3−, (d) Cl−, Br−, I−

FOCUS ON PROBLEM SOLVING EXERCISES 6.87

■ The photoelectric effect can be used in engineering designs for practical applications. For example, infrared goggles used in night-vision applications have materials that give an electrical signal with exposure to the relatively long wavelength IR light. If the energy needed for signal generation is 3.5×10−20 J, what is the minimum wavelength and frequency of light that can be detected?

6.88 Some spacecraft use ion propulsion engines. These engines create thrust by ionizing atoms and then accelerating and expelling them. According to Newton’s laws, this leads to thrust in the opposite direction. Minimizing mass is always important for space applications. Discuss what periodic trends you would need to consider for choosing a material to ionize for engineering designs with this application. 6.89

Laser welding is a technique in which a tightly focused laser beam is used to deposit enough energy to weld metal parts together. Because the entire process can be automated, it is commonly used in many large-scale industries, including the manufacture of automobiles. In order to achieve the desired weld quality, the steel parts being joined must absorb energy at a rate of about 104 W/mm2.

■

■


(Recall that 1 W = 1 J/s.) A particular laser welding system employs a Nd:YAG laser operating at a wavelength of 1.06 mm; at this wavelength steel will absorb about 80% of the incident photons. If the laser beam is focused to illuminate a circular spot with a diameter of 0.02 inch, what is the minimum power (in watts) that the laser must emit per second to reach the 104 W/mm2 threshold? How many photons per second does this correspond to? (For simplicity, assume that the energy from the laser does not penetrate into the metal to any significant depth.) 6.90 Ionization gauges for pressure measurement (see Section 5.7) can be manipulated by changing the voltage of the electron collection grid (see Figure 5.11). Referring to Table 6.4, if the voltage is set to yield electron energies of 1200 kJ/mol, determine what elements the gauge would NOT be able to detect. Given the ability to vary this energy, describe how you would design an apparatus that could distinguish between sodium and magnesium atoms. 6.91 Atomic absorption spectroscopy is based on the atomic spectra of the elements being studied. It can be used to determine the impurities in a metal sample. If an element is present, light at the appropriate wavelength is absorbed. You are working with a metal stamping company and the rolled steel you use to form panels for automobile doors is

■


failing at an alarming rate. There is some chance that the problem is unacceptably high levels of manganese in the steel. Given that the atomic spectrum of manganese has three lines near 403 nm, how could you use a spectrometer to determine the amount of manganese in the steel? 6.92 The red color in fireworks is the result of having strontiumcontaining salts in the fireworks bomb. Similarly, the green/blue colors sometimes seen in fireworks arise from copper salts. Based on your understanding of atomic spectra and the colors in fireworks, describe which atom, copper or strontium, has more widely separated energy levels. Cumulative Problems 6.93 When we say that the existence of atomic spectra tells us that atoms have specific energy levels, are we using deductive or inductive reasoning? 6.94 When Bohr devised his model for the atom, was he using deductive or inductive reasoning? Explain your answer. 6.95 The photochemical reaction that initiates the production of smog involves the decomposition of NO molecules, and the energy needed to break the N!O bond is 1.04 × 10−18 J. (a) Which wavelength of light is needed? (b) How many photons are needed to decompose 0.32 mg of NO?


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7 OUTLINE 7.1

Chemical Bonding and Molecular Structure

INSIGHT INTO Materials for

Biomedical Engineering 7.2

The Ionic Bond

7.3

The Covalent Bond

7.4

Electronegativity and Bond Polarity

7.5

Keeping Track of Bonding: Lewis Structures

7.6

Orbital Overlap and Chemical Bonding

7.7

Hybrid Orbitals

7.8

Shapes of Molecules

7.9

INSIGHT INTO Molecular Scale Engineering For Drug Delivery

TK Biomaterials are designed for use in living systems, so the interactions between such materials and living cells are vital. In this microscope image, the biomaterial is a form of cement used to help repair bones. The smooth structures in the image consist of bone cells that have adhered to the more jagged looking cement. Here this adhesion is the desired effect. But in other cases, strong interactions between cells and biomaterials can have serious negative consequences. H. Xu/American Dental Association Foundation/National Institute of Standards and Technology

T Online homework for this chapter may be assigned in OWL. 200

he study of atomic structure in Chapter 6 provided an interesting detour into modern physics, and there’s no doubt that the paradoxes of quantum mechanics are intellectually intriguing. But quite apart from pure scientific curiosity, the concepts we developed for atomic structure offer important insights into some of the most fundamental questions in chemistry. How and why do atoms combine to form molecules? What do those molecules actually look like, and which factors determine their shapes? In our technological society, the answers to these seemingly abstract questions can have tremendous practical and economic importance. Knowing how to control or selectively exploit the formation of chemical bonds has played a crucial role in the development of many new materials and devices. Many facets of new materials could provide examples of chemical bonding behavior, but here we will highlight the development of materials for use in biomedical applications. Because these materials are designed for use within the human body,

they must conform to important constraints in order to work properly. Virtually all of these constraints are associated with the chemical nature of the materials. In this chapter, we will look at the reasons atoms form bonds with one another, the different types of chemical bonds, and the connections between bonding and molecular shape. An examination of these characteristics can help us understand the properties of many substances and is crucial to understanding how molecules interact and react with one another. We’ll also improve our connection between the microscopic and symbolic views of chemistry as we become able to visualize a molecule’s three-dimensional shape based on its formula.

Chapter Objectives After mastering this chapter, you should be able to ❚ list some factors influencing the biocompatibility of materials and explain how those factors are related to chemical bonding. ❚ use electron configurations to explain why metals tend to form cations whereas nonmetals tend to form anions. ❚ describe the energy changes in the formation of an ionic bond. ❚ define electronegativity and state how electronegativity varies with position in the periodic table. ❚ identify or predict polar, nonpolar, and ionic bonds by comparing electronegativities. ❚ write Lewis electron structures for molecules or ions. ❚ describe chemical bonding using a model based on the overlap of atomic orbitals and recognize some of the limitations of this simple model. ❚ explain how hybridization reconciles observed molecular shapes with the orbital overlap model. ❚ predict the geometry of a molecule from its Lewis structure. ❚ use models (real or software) to help visualize common molecular shapes. ❚ explain the formation of multiple bonds in terms of the overlap of a combination of hybridized and unhybridized atomic orbitals. ❚ identify sigma and pi bonds in a molecule and explain the difference between them.

INSIGHT INTO

7.1

Materials for Biomedical Engineering

From replacement hip joints to heart valves to new lenses for eyes, the collaboration of chemistry and engineering annually provides medical miracles for millions of patients who need to replace worn-out or damaged tissues. Materials used in the human body must be capable of many things, including having physical properties that are at least similar to the biological materials, such as bone, that they are replacing. Just as important, however, is biocompatibility, which is the ability of materials to interact with the natural biological materials without triggering a response from the immune system. Understanding biocompatibility begins with understanding how molecules and materials are put together. For biomedical applications, systems must be engineered to provide the required physical attributes, such as strength and durability, while also allowing the material to be enmeshed with bone or muscle tissues. There are several strategies for achieving biocompatibility. In some cases, the key need is strong cement. The first bone cements used were based on poly(methyl methacrylate) (PMMA). The polymerization reaction, which goes to completion after the cement has been applied, causes the cement to harden. More recently, cements based on calcium phosphate or a composite of calcium phosphate and calcium silicate have been developed. In what

Methyl methacrylate was one of the examples we used to introduce line structures in Chapter 2.

7.1 Materials for Biomedical Engineering

201

ways are these materials similar, and how do they differ in terms of their chemistry? A better understanding of chemical bonding can help us to address these questions. In some cases, materials that could provide a desired physical property may not offer an acceptable level of biocompatibility. The modern engineer, however, need not abandon such materials because it is possible to establish biocompatible coatings. A number of possible methods start with smaller molecules referred to as precursors and react them with each other or with the surface of the object to be coated. Choosing appropriate precursors for coating reactions requires knowledge of chemical bonding not only in these starting materials but also in the resultant coating. The importance of surfaces points to the fact that biomedical engineers must also understand how biomaterials interact with cells in the body. Cell adhesion, like that shown in the chapter opening photo, depends critically on the nature of electrical charges on the surfaces, a property called polarity. Polarity is important not only on the surface of extended systems but also in small molecules in the body, particularly water. We can understand the fundamentals of polarity by looking at the nature of chemical bonding in these systems. The design and implementation of biomaterials clearly involves many complicated factors. Yet the basics of chemical bonding explored in this chapter will allow us to explore many of the questions and challenges posed in this emerging field of engineering.

7.2

The Ionic Bond

All chemical bonds arise from the interaction of oppositely charged particles. The simplest example of such an interaction is found in an ionic bond, formed through the electrostatic attraction of two oppositely charged ions. Many bioengineering applications include ionic substances such as the calcium phosphate or calcium silicate cements that we’ve already mentioned. Ionic compounds can also act as precursors for the growth of ceramic coatings. To understand these systems, we need to examine the forces that dictate which ions will form and how they interact with one another. Binary ionic compounds always form between a metal and a nonmetal, and the more different two elements are in their metallic properties, the more likely it is that a compound formed from these elements will be ionic. So ionic compounds usually consist of elements that are widely separated in the periodic table: a metal from the left hand side and a nonmetal from the right. As we attempt to gain a better understanding of ionic compounds, we’ll want to consider a number of fundamental questions. • In Section 2.5 (page 46), we noted that metallic elements form cations, whereas nonmetals form anions. Why should this be? • How does the energy of attraction between ions compensate for the energy costs of ion formation? • How does the ionic radius affect the strength of the bond between an anion and a cation?

Formation of Cations Recall that ionization energy is the amount of energy required to remove an electron and form a cation. Electron affinity is the amount of energy required to add an electron and form an anion. See Section 6.7, pages 188–191.

202

We’ve already learned that there are energy changes in the formation of ions, as represented by ionization energies and electron affinities. We also know how those energy changes vary as we move through the periodic table. Those trends correlate very well with the location of metals and nonmetals in the table, and this correlation explains the observations that metals form cations and nonmetals form anions. In each case, the ions we find in compounds are those whose formation is not too costly in energy. By looking beyond the first ionization energy, we can also see why some metals form 1+ cations, whereas others have higher charges. As pointed out in our discussion

Chapter 7 Chemical Bonding and Molecular Structure

of Table 6.4 (page 190), there is always a large jump in ionization energy whenever removing an additional electron disrupts an electron configuration ending in np6. This large jump in ionization energy explains whether the cation formed will have a charge of 1+, 2+, etc. For instance, sodium occurs naturally only as singly charged Na+. When we look at the progression in ionization energies for sodium, this is not surprising. The first, second, and third ionization energies for sodium are 496, 4562, and 6912 kJ/mol, respectively. When we see the huge amount of energy required to take a second or third electron from a sodium atom, we can see why nature might choose to stop at only one! Example Problem 7.1 explores the relationship between metal cations and ionization energies.


7.1

Using data from Table 6.4 (page 190), predict the ions that magnesium and aluminum are most likely to form.

Strategy Scan the successive ionization energies for each element to find a point where removing one additional electron causes a dramatic increase in the value. It is not likely that sufficient energy would be available to compensate for such a large increase in ionization energy, so the ion formed will be dictated by the number of electrons lost before that jump in ionization energy.

Solution For magnesium, the first significant jump occurs between the second ionization energy (1451 kJ/mol) and the third (7733 kJ/mol). We would expect two electrons to be removed, and the most commonly found ion should be Mg2+. For aluminum, the third ionization energy is 2745 kJ/mol, whereas the fourth is 11,578 kJ/mol. We would expect three ionizations to occur, so aluminum should form Al3+.

Discussion This type of question is designed to help you see the reasoning behind some of the facts you may already know. In many cases, the way we learn chemistry is by learning information as facts first and then finding out the reason those facts are evident. In the examples from this problem, the answers are consistent with the idea that the ion formed is dictated by the number of electrons lost before the jump in ionization energy. They are also consistent with facts you may already have known, for example, that magnesium will form a 2+ ion because it is an alkaline earth element in Group 2 of the periodic table.

Check Your Understanding The calcium in calcium phosphate exists as Ca2+ ions. Based on this fact, what would be your qualitative prediction for the values of the successive ionization potentials for calcium?

As our examples show, the most common cation for metals in the s or p blocks of the periodic table is generally one with an np6 configuration. But this idea does not carry over to transition metals. As we noted in Chapter 2, many transition metals form more than one stable cation. Transition metals have partially filled d subshells, but ionization usually begins with loss of electrons from the s subshell. When iron forms Fe2+, for example, the electron configuration changes from 1s 22s 22p63s 23p64s 23d 6 in the neutral atom to 1s 22s 22p63s 23p63d 6 in the ion. The loss of one additional electron from a d orbital leaves Fe3+ with a half-filled d subshell. This half-filled subshell is a fairly stable arrangement, so the Fe2+ and Fe3+ ions are both stable. With many transition metals, the situation can become quite complicated, and we will not explore this in greater detail.

The idea of s, p, d, and f blocks in the periodic table was introduced in Figure 6.19, page 186.

If a shell is half-filled, no electrons are paired in any orbitals. Pairing of electrons is allowed (with opposite spins), but it still increases electron– electron repulsions.

7.2

The Ionic Bond

203

Formation of Anions If we look at the periodic trends for electron affinity (Figure 6.22, page 191) we see that they are similar to those for ionization energies. Recall that electron affinity is the energy change for the formation of an anion from a neutral atom: X(g) + e– : X–(g)

1 picometer (pm) = 10−12 m

This process is often energetically favored, which means that energy is released when the anion is formed. The amount of energy released when atoms acquire additional electrons tends to increase as we go from left to right across the periodic table. Looking at this trend, those atoms that gain electrons most readily also would need to lose large numbers of electrons to produce cations with stable np6 electron configurations. Thus nonmetals tend to form anions rather than cations. For the formation of anions from nonmetals, we do not find any of the complexities we encountered with the transition metals because all of the nonmetals belong to the main groups and do not have partially filled d subshells. They behave like the main group metals, in that they tend to acquire electron configurations with filled p subshells. The primary difference is that they do so by gaining, rather than losing, electrons. Thus, all the halogens form singly charged (1–) anions, resulting in np6 electron configurations. Similarly, oxygen and sulfur form anions with a 2– charge, resulting in the same np6 configuration. In addition to charge, the size of an ion ultimately has an effect on the strength of an ionic bond, largely due to geometric constraints. In Section 6.7, we noted that the atomic radius increases going down a group and decreases going across a period. What trend should we expect for the radii of cations and anions? The loss of an electron to form a cation decreases electron–electron repulsion, leaving the remaining electrons more tightly bound to the nucleus. So we would expect positive ions to be smaller than their respective parent atoms. On the other hand, adding an electron to form an anion will lead to increased repulsion among the valence electrons. So anions should be larger than their corresponding neutral atoms. Ionic radii can be determined very accurately by measuring the spacing in crystal lattices, and the data are consistent with these expectations. With an ionic radius of 102 pm, Na+ is significantly smaller than the sodium atom, which has an atomic radius of 186 pm. And the 133-pm ionic radius of F – is significantly larger than the 72-pm atomic radius of neutral fluorine. Figure 7.1 shows atomic and ionic radii, and the observed trends are clear. When we compare similarly charged ions going down a group, ionic size increases in the same way that we saw for atomic size. For cations, the trend in ionic size going across a period also mimics the trend we saw for neutral atom—decreasing as we go from left to right. This is easily understood. Considering Na+ and Mg2+, for example, we see that both ions have the same electron configuration. But Mg2+ has one additional proton in its nucleus, and so it will hold its electrons more tightly, leading to a smaller ion. We can apply the same sort of reasoning to the common anions in the p block. The oxide ion, O2–, and the fluoride ion, F–, have the same electron configuration. But fluoride has one additional proton in its nucleus, and so it is the smaller ion. Similarly, Cl– is slightly smaller than S2–. The formation of isolated ions, of course, does not constitute an ionic bond. Indeed, starting with a neutral metal atom and a neutral nonmetal atom, it will usually cost energy to form an ion pair (i.e., a cation and an anion). Ionization energies are always positive, and the energy input to produce a cation from a metal is not offset by the energy released when an anion is formed from a nonmetal. But once formed, the oppositely charged ions will attract one another, leading to a significant lowering of their overall energy. As we saw in Section 2.3, Coulomb’s law says that the force between two charges is proportional to the product of the magnitude of their charges divided by the square of the distance between them. q1q2 F ∝ —— r2

204


(7.1)

Group 1 Ion

Group 2 Atom

Ion

Group 13

Atom

Ion

Group 16 Atom

Ion

Group 17 Atom Ion

Atom

Li+ 76

Li 152

Be2+ 45

Be 112

B3+ 23

B 85

O2– 140

O 73

F– 133

F 72

Na+ 102

Na 186

Mg2+ 66

Mg 160

Al3+ 51

Al 143

S2– 184

S 103

Cl– 181

Cl 100

K+ 151

K 227

Ca2+ 100

Ca 197

Ga3+ 62

Ga 135

Se2– 198

Se 119

Br– 196

Br 114

Rb+ 161

Rb 248

Sr2+ 126

Sr 215

In3+ 80

In 167

Te2– 207

Te 142

I– 220

I 133

Figure 7.1 ❚ The nested spheres in the figure illustrate how the sizes of the indicated ions compare to the corresponding neutral atoms. Metals and their cations are shown on the left and nonmetals and their anions on the right. All radii are given in picometers.

Here, F is the force, q1 and q2 are the charges, and r is the distance between them (in this case, the distance between the nuclei of the two ions). It can also be shown that the potential energy, V, of the system is given by q1q2 V = k —— r

(7.2)

Here, k is a proportionality constant, whose value is 1.389 × 105 kJ pm/mol if the charges are expressed in units of electron charge. Let’s look at a specific case by evaluating the energy change in the formation of a bound NaF ion pair from isolated atoms of sodium and fluorine. The ionization energy of sodium is 496 kJ/mol, and the electron affinity of fluorine is –328 kJ/mol. So we need to supply 496 kJ/mol to ionize the sodium, but 328 kJ/mol will be released in the formation of the fluoride ion. Combining these two numbers, we see that forming the individual ions will cost 168 kJ/mol (496 − 328 = 168). Next, we need to use Equation 7.2 to determine the energy released by the coulombic attraction between the anion and cation. The radii of Na+ and F – are 102 pm and 133 pm, respectively, so the internuclear distance should be 235 pm. Inserting this into Equation 7.2 as r and evaluating gives us the energy of attraction between the ions: kJ pm (+1)(−1) V = q1.389 × 105 ———r ————— = −591 kJ/mol mol (235 pm) The negative sign indicates that this energy is released. Notice that the magnitude of the coulombic energy far outweighs the energetic cost of forming the two ions, so the whole process we have described does release energy. Before making too much of that conclusion, though, we should consider two ways in which our calculation is not realistic. First, we began with isolated atoms of sodium and fluorine. If we were to react elemental sodium and fluorine, though, we would not start with isolated gas phase atoms. Under any reasonable conditions, sodium would be a metallic solid and fluorine would be a diatomic gas. So there would be some added energy input needed to produce the individual atoms with which we started. On the other end of our calculation, though, we found the coulombic energy for a

7.2

The Ionic Bond

205

Black lines show attractive forces between (1st) nearest neighbors.

⫹ ⫺

Red lines show repulsive forces with 2nd nearest neighbors. Blue lines show attractive forces with 3rd nearest neighbors. Fourth nearest neighbor interactions and higher are not indicated in the figure.

⫹ ⫺

⫺⭈ ⫹ ⫺⭈ ⫹

⫹

⫺⭈ ⫹

⫹ ⫺⭈ ⫹ ⫺⭈ ⫹

⫺⭈ ⫹ ⫺⭈ ⫹

⫺ ⫺⭈

⫺⭈ ⫹

⫹ ⫺⭈ ⫹ ⫺⭈ ⫹

⫺⭈

Figure 7.2 ❚ In a solid lattice, any given ion has a large number of interactions with other ions, some attractive and some repulsive. This two-dimensional representation of a lattice shows the interactions of the nearest neighbors, second nearest neighbors, and third nearest neighbors.

The overall result of all attractive and repulsive forces in an ionic crystal is estimated by calculating a Madelung constant, a mathematical series that accounts for all of the interactions in a particular crystal geometry.

single ion pair. But if we really did form sodium fluoride, we would produce a bulk crystal lattice made up of many ions. In such an ionic lattice, ions are attracted to every other oppositely charged ion. At the same time, ions of like charge are repelling each other. The farther one ion is from another, the weaker the attraction or repulsion (Figure 7.2). Ultimately, the lattice energy of a crystal contains contributions from all of these attractions and repulsions. In the NaF structure, each sodium ion is surrounded by six fluoride ions, and each fluoride has six neighboring sodium atoms. To get an accurate indication of the coulombic energy of the lattice, we would need to account for the interaction of each ion with its many neighbors. So, although starting with isolated atoms leads to an overestimation of the energy released in forming the ionic bond, using a single ion pair instead of an extended lattice greatly underestimates the energy. The actual energy released in the reaction between sodium metal and diatomic fluorine gas to produce solid sodium fluoride turns out to be 577 kJ/mol. From the form of Equation 7.2, we can see that small, highly charged ions will tend to form ionic compounds with large lattice energies, and large ions with small charges will tend to form compounds with small lattice energies. Example Problem 7.2 explores this connection.

E XAM PL E PRO BLEM

7.2

In each of the following pairs of ionic solids, the ions are arranged into similar lattices. Which one has the largest lattice energy: (a) CaO or KF? (b) NaF or CsI?

Strategy Questions like this require us to find some basis for comparison. In this particular case, the factors associated with lattice energy have to do with the ions that make up the compound. We have two factors to consider. Smaller ions tend to have stronger coulombic interactions because the distance between the nuclei is smaller, and ions with larger charges have stronger interactions. An examination of these two factors may allow us to predict which substance has the larger lattice energy.

Solution (a) In CaO, calcium exists as Ca2+ and oxygen as the oxide ion, O2–. In KF, the constituent ions are K+ and F–. Because both are from elements in the same row of the periodic table, the sizes of the two cations should not be too different. The same is true for the anions, so to a first approximation, we might discount the size factor. CaO should have the larger lattice energy because the charges on its ions are larger than the charges on the ions in KF. (Note that if we wanted to look more closely at the sizes, we would expect that K+ would be slightly larger than Ca2+ and O2– would

206


be slightly larger than F–. Because the larger cation is in KF and the larger anion is in CaO, our assumption that size would not be very important seems good.) (b) All of the ions in this problem are singly charged, so the size difference must be the key issue. Na+ and F– are smaller than Cs+ and I–, so the lattice spacing in NaF must be smaller. Therefore, NaF should have the larger lattice energy.

Discussion To get started here, we needed to determine the identities of the ions involved. One way to do this is to consider the jump in ionization energy we explored in Example Problem 7.1. However, with just a little practice, it is easy to determine the likely charges for most main group elements based on their location in the periodic table. Ionic charge and radius were taken into consideration, based on the positions of the ions in the periodic table. The use of this “periodic reasoning” points out the importance of learning the trends in the periodic table. Note, however, that to be able to make the comparisons in this problem, the lattice structures of the compounds in question must be similar.

Check Your Understanding Of the three compounds, CaF2, KCl, and RbBr, which should we predict to have the smallest lattice energy? Why?

7.3

The Covalent Bond

Consideration of ionization energy, electron affinity, and coulombic potential is fine for ionic compounds. But what about covalent bonds like those found in polymers? We noted back in Section 2.4 that covalent bonds are based on the sharing of pairs of electrons between two atoms. How does covalent bonding occur, and how does it influence the behavior of compounds like the poly(methyl methacrylate) used in bone cements?

Chemical Bonds and Energy When chemical bonds are formed between nonmetal elements, valence electrons are shared rather than transferred. But just as for ions, we must expect that the driving force behind bond formation should be a lowering of the overall energy. The energetic advantage of forming a covalent bond can be illustrated by considering how the potential energy would vary in a thought experiment in which two atoms approach each other. If the atoms are sufficiently far apart, they won’t exert any discernible force on one another. (You might say that neither one knows the other is there.) But as the atoms begin to come closer together, eventually the electrons from one atom can become attracted to the positively charged nucleus of the other atom. This added attraction lowers the potential energy of the system. Of course, the two nuclei will repel each other, as do the bonding electrons, so the nuclei cannot approach too close to one another. At some point, a balance between these attractions and repulsions will occur, and the energy of the system is minimized (Figure 7.3). This point of minimum energy corresponds to the formation of a covalent bond. The energy released when isolated atoms form a covalent bond is called the bond energy, and the distance between the nuclei of the bonded atoms is called the bond length. When electrons are shared between two atoms, their distribution is different from that in the individual atoms. In Figure 7.4, we see how electron density in a molecule differs from the spherically symmetrical distribution seen in isolated atoms. In a covalently bonded molecule, there is a buildup of electron density, or a concentration of negative charge, in the region between the two bonded atoms. The increase

It is important to remember that bond formation ALWAYS releases energy.

7.3

The Covalent Bond

207

4

When the atoms are pushed too close to each other, the nuclei start to repel each other. There is a dramatic increase in the potential energy as a result.

1 When widely separated, atoms do not interact with each other at all, so potential energy is zero. Increasing energy

Figure 7.3 ❚ A schematic diagram shows how the potential energy of a pair of atoms might vary as the atoms draw closer together and form a covalent bond. The minimum in this “potential well” corresponds to the normal bond length and bond energy for the diatomic molecule.

0

2 Increasing internuclear distance The amount by which the energy is lowered by this interaction is called the bond energy.

Figure 7.4 ❚ When two atoms approach one another, the negatively charged electron clouds are each attracted to the other atom’s positively charged nucleus. This attractive interaction potential is largely responsible for the reduction of energy in the formation of a chemical bond.

As atoms approach each other they begin to interact and the potential energy decreases. Attractive interactions lower the potenial energy.

3 When the atoms reach an optimal separation, there is significant sharing of electrons between the two nuclei. This distance lowers the potential energy more than any other and is called the bond distance.

When atoms are widely separated, their electron clouds are not distorted by the presence of the other atom and are spherical.

As atoms approach, the negatively charged electron clouds become attracted to both nuclei.

When atoms are close enough to form a chemical bond, significant electron density builds up between the nuclei and the electron clouds are no longer spherical.

in electron density need not be particularly large; even a slight increase of negative charge in this region lowers the potential energy of the system and can stabilize the molecule. The formation of a chemical bond always releases energy. Once a bond is formed, that same amount of energy—the bond energy—must be supplied to break the bond

208


apart. The magnitude of this bond energy depends on the atoms involved. For example, when two moles of fluorine atoms combine to form one mole of molecular fluorine, 156 kJ of energy is released. So we say that the F—F bond energy is 156 kJ per mole. Compared to many other molecules, F2 is held together by a relatively weak covalent bond. Hydrogen fluoride, for example, has a bond energy of 565 kJ/mol. This means that its bond is more than three times stronger. Because more energy is required to break its atoms apart, we say that the bond in HF is more stable than the bond in F2. When we want to assess the stability of a compound, some general ideas about the different types of bonds present are useful.

Chemical Bonds and Reactions Chemical reactions involve rearranging bonds, so that reactants are transformed into products. A reaction will be energetically favored if the energy required to break bonds is less than that released in making new bonds. Some biomaterials such as a Teflon® mesh used to reinforce an abdominal wall after surgery are intended to replace or strengthen tissues. Clearly it is important that these materials not be consumed by metabolic reactions in the body. Teflon is a trade name for polytetrafluoroethylene (PTFE), and a key property of PTFE is its resistance to chemical reactions. Metabolism is fairly similar to combustion. So many biomolecules are broken down into H2O and CO2. We can examine the resistance of PTFE to metabolism by comparing the combustion of PTFE with that of the hydrocarbon component of biomolecules. PTFE is a fluorocarbon, which you might think of as a hydrocarbon in which all of the hydrogen atoms have been replaced by fluorine. So the combustion of Teflon would have to produce CO2 and OF2. Table 7.1 shows the bond energies of various types of bonds that are relevant to this discussion. Note that a C!F bond is slightly stronger than a C!H bond and thus is harder to break. More importantly, notice the large difference in strength between O!H and O!F bonds. These bond energies show us that to burn a fluorocarbon, we would have to replace strong C!F bonds with weaker O!F bonds. Energetically, this is not at all favorable, and many fluorinecontaining molecules are not flammable. Thus a simple look at bond energies can help explain the nonreactive nature of Teflon.

The familiar use of Teflon in cooking products shows that PTFE tends not to react even at relatively high temperatures.

Chemical Bonds and the Structure of Molecules To explain chemical bonding, we must go beyond the discussion of energetics in the previous section. When we looked at ionic bonding between metals and nonmetals, we saw that the formation of cations and anions is consistent with achieving np6 electron configurations that resemble those of the rare gases. For nonmetal atoms, a rare gas configuration can be achieved by gaining the electrons from metal atoms, leading to ionic bonds. When nonmetals interact with other nonmetals, however, they share electrons, effectively giving each atom access to all the shared (bonding) electrons. Thus, by sharing a suitable Table

❚

7.1

Bond energies of some types of bonds important in the combustion of hydrocarbons and fluorocarbons

Type of Bond

Bond Energy (kJ/mol)

C!F

485

C!H

415

O!F

190

O!H

460

7.3

The Covalent Bond

209

••

•

••

Ne ••

••

•

•

•

•

••

F ••

••

••

O ••

•

•

•

••

N

•

•

•

C

•

B

•

•

Be

•

Li

•

Lewis dot symbols of atoms help explain the valence of elements. The original definition of valence was derived from the number of chemical bonds that an element would form. Counting unpaired electrons in the Lewis symbol provides the valence of elements in the s and p blocks of the periodic table.

number of electrons, an atom can achieve an electron configuration that resembles that of a rare gas. For many molecules composed of elements from the main groups, this feature leads to eight valence electrons around each atom. This observation is the basis for what is known as the octet rule: An atom will form covalent bonds to achieve a complement of eight valence electrons. Recall that a rare gas atom has a valence shell electron configuration of ns2np6, for a total of eight valence electrons. Because it has no p orbitals in its n = 1 valence shell, hydrogen is an exception to the octet rule and can share only two electrons. There are other exceptions as well, which we will note as we encounter them. The basic concept of the octet rule can be conveniently depicted using a Lewis dot symbol, as developed by the American chemist G. N. Lewis in the early 20th century. These Lewis symbols help keep track of valence electrons, making it easier to predict bonding in molecules. They are particularly useful for compounds of main group elements. To draw a Lewis symbol for an element, we use its atomic symbol and place dots around it on four sides to depict the number of valence electrons. As we go across a period in the main groups, the first four dots are placed singly around the symbol. Starting with the fifth dot, they are paired. Thus the elements in the second period would have the following Lewis symbols:

•

• •

•

•

•

•

•

Ge

•

Si

•

C

•

•

Because elements in the same group have the same number of valence electrons, the Lewis symbols for subsequent periods are the same, except for the identity of the elements. For example, the Lewis dot symbols for carbon, silicon, and germanium are as follows:

Lewis symbols for the main group elements are shown in Figure 7.5. These symbols for individual atoms are of fairly little use, but similar representations for molecules are much more powerful. When we draw Lewis dot representations of molecules, we add one more level of symbolism to denote bonds between atoms. A line between two atomic symbols represents a shared pair of electrons forming a bond; thus when counting electrons, one line is equivalent to two dots. Figure 7.6 shows the Lewis symbol for the hydrogen atom and the Lewis structure of the simplest molecule, H2. A Lewis structure shows how electrons are shared in a molecule. The Lewis structure for F2 is shown in Figure 7.7. The sharing of electrons to form a bond Figure 7.5 ❚ The Lewis symbols for elements in the first three periods of the periodic table are shown. Notice the similarities among elements in the same group. Lewis structures help keep track of electrons involved in chemical bonds.

210

Group

1

2

13

14

15

16

17

18

7

8 (except He)

Number of electrons in valence shell

1

Period 1

H

Period 2

Li

Be

B

C

N

O

F

Ne

Period 3

Na

Mg

Al

Si

P

S

Cl

Ar

Period 4

K

Ca

Ga

Ge

As

Se

Br

Kr

Period 5

Rb

Sr

In

Sn

Sb

Te

I

Xe

Period 6

Cs

Ba

Tl

Pb

Bi

Po

At

Rn

Period 7

Fr

Ra


2

3

4

5

6

He

H• ⫹ •H

H •• H or H

H

Figure 7.6 ❚ By sharing the electron from each atom, two hydrogen atoms can form a chemical

or

••

F ••

••

F ••

••

••

••

••

••

••

F F •• •• ••

••

••

••

• ⫹ •F F •• ••

••

bond. In doing so, each atom has enough valence electrons around it to achieve a filled n = 1 shell.

Figure 7.7 ❚ When two fluorine atoms combine, they form a stable chemical bond. The Lewis structures for the F atoms show why just one bond forms. By sharing one pair of electrons, each atom is surrounded by eight valence electrons.

effectively gives each fluorine atom a full octet of electrons. When fluorine is part of a covalently bonded system, it virtually always forms a single bond by sharing an electron pair, thereby satisfying the octet rule. A pair of electrons shared by two atoms is referred to as a bonding pair. The other paired electrons that are associated with a single atom are referred to as nonbonding electrons, or as lone pairs. A fluorine molecule has one bonding pair and six lone pairs. Of course, most molecules are more complicated than H2 or F2 and will contain more than one pair of shared electrons. If two oxygen atoms share a single pair of electrons, for example, then neither atom has an octet of valence electrons. However, if two pairs of electrons are shared, then both oxygen atoms conform to the octet rule. For diatomic nitrogen, three pairs of electrons must be shared to satisfy the octet rule. Multiple bonding results from the sharing of more than one electron pair. When two pairs of electrons are shared, the result is a double bond, and three shared pairs result in a triple bond. Double bonds are stronger than single bonds, and triple bonds are stronger than double bonds. However, a double bond, although stronger than a single bond, is generally not twice as strong. The bond energies below for carbon–carbon bonds provide a good example.

Type of Bond

Bond Energy (kJ /mol)

C!C

346

C"C

602

C#C

835

The relative strengths of multiple bonds have important consequences. The polymerization of methyl methacrylate to form PMMA provides a good example. The methyl methacrylate monomer contains a carbon–carbon double bond. In the polymerization reaction, two of the electrons forming this double bond are used to create a bond to a second monomer unit. The end result is that the double bond in the reactant is transformed into two single bonds in the product. Because the energy released in forming two single bonds is greater than the energy needed to break a double bond, the reaction to form PMMA is an energetically favorable process. This is true for many similar reactions in which a reactant with a double bond is converted into a product with two single bonds. It explains why the common reaction of the C"C functional group is addition. The same is true for triple bonds, although those are much less common.

7.4

This polymerization reaction is analogous to the formation of polyethylene, which was discussed in Section 2.8.

Electronegativity and Bond Polarity

Having considered both ionic and covalent bonding, we can begin to understand the chemical properties of both polymers and ceramics. One key biomaterial is hydroxyapatite, Ca10(PO4)6(OH)2, which involves both ionic and covalent bonding. How can we predict which bonding interactions will be ionic and which will be covalent?

7.4 Electronegativity and Bond Polarity

211

The sharing of electrons to form covalent bonds and the outright transfer of electrons that occurs in ionic bonding represent two ends of a bonding continuum. To describe compounds whose bonding falls somewhere in the middle of this continuum, we will introduce two additional concepts: electronegativity and bond polarity.

Electronegativity We’ve seen that ionic bonds form between a metal and a nonmetal, with one or more electrons being transferred from the metal to the nonmetal. Compared to metals, nonmetals are relatively “greedy” for electrons. Even in covalent bonds, the small size and high electron affinity of many nonmetals mean that electrons tend to be strongly attracted to those atoms. The attraction of an atom for the shared electrons in a covalent bond is called electronegativity. Although this concept may seem simple enough, quantifying it is very difficult. Unlike ionization energy or electron affinity, electronegativity is not an energy change that can be determined experimentally. Instead, electronegativity values are established by considering a number of factors, such as atomic size, electron affinity, and ionization energy, and a scale of arbitrary units has been chosen. Figure 7.8 gives the resulting values for the electronegativities of most elements. The higher the electronegativity value, the more likely an element is to attract extra electron density when it forms chemical compounds. Electronegativity tends to increase as we go up a group (decreasing atomic number) or across a period from left to right. Thus the trends in electronegativity parallel those in electron affinity. Fluorine, the most electronegative element, is assigned a value of 4.0. Oxygen and nitrogen also have relatively high electronegativities. When these elements take part in covalent bonding, they tend to take “more than their fair share” of the bonding electron pair. What are the implications for a molecule when electrons are shared unequally?

VIIIA

IA Metals Nonmetals Metalloids

1

1

H 2.1

2

3

4

5

6

7

IIA

2

IIIA

IVA

VA

VIA VIIA

He

3

4

5

6

7

8

9

10

Li

Be

B

C

N

O

F

Ne

1.0

1.5

2.0

2.5

3.0

3.5

4.0

11

12

13

14

15

16

17

18

Na

Mg

VB

IB

IIB

P

S

Cl

Ar

1.2

IVB

Si

1.0

IIIB

Al 1.5

1.8

2.1

2.5

3.0

19

20

21

22

23

24

25

26

27

28

29

30

31

32

33

34

35

36

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

Kr

0.9

1.0

1.3

1.4

1.5

1.6

1.6

1.7

1.7

1.8

1.8

1.6

1.7

1.9

2.1

2.4

2.8

37

38

39

40

41

42

43

44

45

46

47

48

49

50

51

52

53

54

Rb

Sr

Y

Zr

Nb

Mo

Tc

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

I

Xe

0.9

1.0

1.2

1.3

1.5

1.6

1.7

1.8

1.8

1.8

1.6

1.6

1.6

1.8

1.9

2.1

2.5

55

56

57

72

73

74

75

76

77

78

79

80

81

82

83

84

85

86

Cs

Ba

La

Hf

Ta

W

Re

Os

Ir

Pt

Au

Hg

Tl

Pb

Bi

Po

At

Rn

0.8

1.0

1.1

1.3

1.4

1.5

1.7

1.9

1.9

1.8

1.9

1.7

1.6

1.7

1.8

1.9

2.1

87

88

89

Fr

Ra

Ac

0.8

1.0

1.1

*

VIIIB

VIB VIIB

†

*

†

58

59

60

61

62

63

64

65

66

67

68

69

70

71

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.1

1.0

1.2

90

91

92

93

94

95

96

97

98

99

100

101

102

103

Th

Pa

U

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

1.2

1.3

1.5

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.5

Figure 7.8 ❚ Electronegativity values for the elements are shown. Notice how the values increase toward the right and top of the periodic table. 212


When separated, both hydrogen and fluorine are spherical. The negative charge of the electrons and the positive charge of the nucleus offset each other.

When bonded, the more electronegative fluorine attracts the shared electrons more than hydrogen. The electron density shifts causing a partial separation of charge. δ–

δ+

H

Figure 7.9 ❚ The highly electronegative fluorine atom attracts electrons more strongly than hydrogen does. So the two form a polar covalent bond.

H—F

F

Bond Polarity The high electronegativity of fluorine plays an important role in a number of processes. The fluoridation of teeth, for example, is one of the earliest manipulations of biomaterials and involves the replacement of an !OH group in hydroxyapatite to form fluorapatite. To understand how fluorine participates in chemical bonds from covalent to ionic, a convenient starting point is to compare the electron distribution around the atoms in a molecule with the electron distribution around the isolated atoms. If we observe a fluorine atom from a distance, it appears to be electrically neutral because the nuclear and electronic charges offset each other. Because fluorine is the most electronegative element, when a fluorine atom participates in a covalent bond with another element, such as hydrogen, more than half of the electron distribution of the bonding electron pair is associated with the fluorine atom. The result is that the fluorine atom becomes a center of partial negative charge because it has more electron density than is necessary to balance its own nuclear charge. At the same time, the hydrogen atom at the opposite end of the bond takes on a partial positive charge because it has insufficient electron density to balance its nuclear charge (Figure 7.9). Thus, covalent bonding between atoms of different electronegativity results in a bond in which there is a separation of positive and negative charge centers. Because the molecule has one region of positive charge and another of negative charge, the molecule has an electric field associated with it. These two points of positive and negative charge constitute a dipole, and such a bond is referred to as a polar bond. Because the electrons are still shared rather than transferred, these bonds are often called polar covalent bonds. Polar covalent bonding completes the continuum of chemical bonding. The polarity of covalent bonds varies from completely nonpolar (as in F2 and H2), through varying intermediate degrees of polarity (as in HF), to completely ionic (as in NaF). The type of bond formed depends on the electronegativity difference between the bonding atoms. The greater the difference in electronegativity, the more polar the bond. At some point, typically when the difference in electronegativity exceeds about 2.0, the bond is classified as ionic. Using the electronegativity values listed in Figure 7.8, we see that NaF, with an electronegativity difference of 3.0, is clearly ionic. The C!F bond in a fluorocarbon such as Teflon is polar covalent (electronegativity difference = 1.5), and the F!F bond is completely covalent (electronegativity difference = 0). Example Problem 7.3 explores this concept further.


The relationship between the difference in electronegativity and the ionic character of a bond is actually continuous, but using a cutoff to label a bond as ionic or polar simply makes the language more tractable.

7. 3

Which bond is the most polar: C!H, O!H, or H!Cl? For each of these bonds, determine which atom has a partial positive charge.

Strategy Bond polarity depends on the difference in electronegativity between the two bonded atoms. So we can look up the electronegativity values and subtract to find the difference in each case. Negatively charged electrons are drawn more strongly 7.4 Electronegativity and Bond Polarity

213

toward an atom with higher electronegativity, so the less electronegative atom has a partial positive charge.

Solution From Figure 7.8, we find the following electronegativities: H = 2.1,

C = 2.5, O = 3.5, Cl = 3.0. The electronegativity differences are, therefore, C!H = 0.4, O!H = 1.4, and H!Cl = 0.9. So the O!H bond is the most polar of the three. In each of the three bonds considered here, H is the less electronegative atom, so it carries the positive charge.

Discussion This question may appear simple because we are using electronegativity values and differences in a straightforward way. Despite this simplicity, the results can provide important insight into material properties. Our focus at this point is on single molecules, but invariably in laboratory observations in the real world, we encounter huge collections of molecules. The way these samples of molecules behave is often influenced strongly by the polarity of the bonding within the individual molecules. Therefore, understanding a bond dipole and predicting its presence using electronegativity will be useful in a number of ways as we move forward in our study of chemistry.

Check Your Understanding Which bond should be nonpolar: C!N, N!Cl, or H!Br?

The term used to describe the breaking open of red blood cells is hemolysis.

Bond polarity plays an important role in several aspects of biomaterials, particularly with regard to biocompatibility. Materials that will be deployed in living system always interact with those systems. A key example is the interaction of biomaterials with blood. Polar bonds in molecules along the outer wall of blood cells help to allow the blood to interact with water. Those same polar bonds can also interact with biomaterials and influence biocompatibility. Amorphous silica is composed of a large number of silica units, SiO4. Most sand is silica, and we don’t usually think of sand as toxic, but amorphous silica is not biocompatible. In particular, when this material was tested for compatibility, it was found to destroy blood cells, essentially allowing the hemoglobin that transports oxygen to leak out. An explanation for this behavior can be found in the fact that the surface of amorphous silica has a large number of polar Si!O bonds. These bonds are capable of interacting strongly with positive charges on the surface of a red blood cell, as illustrated in Figure 7.10. The interaction is so strong that the silica particle remains close to the blood cell through mechanical jostling and ultimately pierces the cell membrane, destroying the cell.

Figure 7.10 ❚ Cells often have polar bonds on their exterior because they interact with polar water molecules. A substance like amorphous silica (left) that also has polar surface species can interact strongly with the polar bonds on the surface of a cell such as the blood cell depicted on the right. These strong interactions can lead to cell damage, causing biocompatibility concerns.

⫺

O

Si

⫺

+N

O

⫺

O O⫺

O⫺

Si

O⫺ O⫺ ⫺ O Si O⫺ O⫺ ⫺ O Si O⫺ O⫺ ⫺ O Si O⫺ O⫺

+N +N +N +N +N

Si O ⫺ O⫺

214


+N +N

7.5


We’ve now surveyed some important aspects of chemical bonding, and we’ve gained an understanding of a number of key concepts. But if we want to think about the connections between bonding and the properties of compounds, we often need to know not just that bonds have formed, but how many bonds have formed and between which elements. The Lewis structures we introduced in Section 7.3 provide a powerful tool for doing this. In drawing Lewis structures for very simple molecules like H2 and F2, we drew each atom separately and then brought them together to form a bond. This approach obviously assumes that we know which atoms are to be attached to one another. The choice is obvious when we are dealing with a diatomic molecule but can become problematic for larger molecules. If we want to be able to draw Lewis structures for more complicated molecules or ions, it is useful to use an algorithmic approach. By applying such an algorithm carefully, we can predict which bonds are likely to be formed by examining and comparing the possible structures for a specific molecule. As we introduce our algorithm for drawing Lewis structures, it will help if we use a simple molecule to illustrate each step. We could choose water, but the fact that hydrogen does not follow the octet rule means that is not the best example. Instead, we will use a very similar molecule, oxygen difluoride (OF2) as our first example. Step 1: Count the total number of valence electrons in the molecule or ion. We can do this by summing the numbers of valence electrons for each atom in a molecule. (If we were trying to draw the structure for a polyatomic ion, we would also need to account for the ion’s charge, as we’ll see in Example Problem 7.5.) For OF2, each fluorine atom has seven valence electrons, and the oxygen atom has six. So we will have a total of 20 valence electrons in the molecule. F

2 × 7 = 14

O

1×6=6

Although it may seem routine after a little practice, it is always vitally important to carry out Step 1 of this process. Using the wrong number of valence electrons is among the most common errors that students make in drawing Lewis structures.

Total = 20 Step 2: Draw the skeletal structure of the molecule. Here, we can rely on some general rules. The element written first in the formula is usually the central atom in the molecule, except when hydrogen is listed first. (Hydrogen can never be a central atom because it can form only one bond!) Usually, the central atom is the least electronegative. Sometimes more than one possible skeleton can be drawn. In such cases, we might rely on chemical intuition to choose likely alternatives. Or we could also draw multiple possible structures and evaluate those to see which seemed more plausible. In OF2, oxygen is written first and is the less electronegative element, so it is the central atom. We write the two fluorine atoms surrounding the oxygen. (It isn’t important at this point on which two sides of the oxygen symbol we choose to place the fluorine atoms.) F

O F

Step 3: Place single bonds between all connected atoms in the structure by drawing lines between them. We don’t yet know how many bonds might be needed in our molecule, but we do know that at least one bond is needed to attach each of the outer fluorine atoms to the central oxygen atom, to keep the molecule from falling apart. So we can go ahead and draw two single bonds. F

O F

7.5


215

Step 4: Place the remaining valence electrons not accounted for in Step 3 on individual atoms until the octet rule is satisfied. Place electrons as lone pairs wherever possible. It is usually best to begin placing electrons on the outer atoms first, leaving the central atom(s) for last. Remember that each bond contains two electrons, so in drawing two single bonds, we have accounted for four electrons. Our molecule has 20 valence electrons, so 16 remain to be added to our structure. Placing six of these electrons around each fluorine atom brings us up to 16, and placing the remaining four on the central oxygen atom accounts for all 20 available electrons. All 16 of these electrons should be drawn as lone pairs. F

O F

At this point, we know that all of the available valence electrons appear somewhere in our structure. To decide whether this is the optimal arrangement of the valence electrons, we check to see if each atom satisfies the octet rule. For OF2, our structure now has a full octet of electrons on each atom, so this is the final structure. In other cases, we might find one or more atoms that have less than a full octet; then, we would take an additional step. Step 5: Create multiple bonds by shifting lone pairs into bonding positions as needed for any atoms that do not have a full octet of valence electrons. Our structure for OF2 already satisfies the octet rule, so multiple bonds are not required in this molecule. We will work with this added step in future structures. The following examples will further illustrate our algorithm. Of course, the best way to become familiar with these steps is simply to practice drawing Lewis structures yourself.

E XAM PL E PRO BLEM

7.4

Chlorofluorocarbons have been used as refrigerants since the 1930s, but their use has largely been phased out worldwide because of concerns about atmospheric ozone depletion. Draw the Lewis structure of dichlorodifluoromethane, CF2Cl2, also known as DuPont’s Freon-12.

Strategy Use the algorithm for drawing Lewis structures. Solution Step 1: Count the total number of valence electrons. Carbon has four, fluorine has seven, and chlorine has seven. C

1×4= 4

F

2 × 7 = 14

Cl

2 × 7 = 14 Total = 32

Step 2: Draw the skeletal structure of the molecule. Carbon is listed first in the formula, and it is the least electronegative, so use it as the central atom. Cl F

C Cl

216


F

Step 3: Place single bonds between all connected atoms. Cl F

C

F

Cl Step 4: Place the remaining valence electrons not accounted for in Step 3 on individual atoms until the octet rule is satisfied. Four single bonds have been drawn, and they require eight valence electrons. We have a total of 32 available electrons, so we are left with 32 – 8 = 24. This gives each of the four halogen atoms three lone pairs (6 electrons). Cl F

C

F

Cl Step 5: Create multiple bonds for any atoms that do not have a full octet of valence electrons. Check to make sure each atom has an octet of eight electrons. Carbon has four bonding pairs, or eight electrons. Each halogen has one bonding pair and three lone pairs, so they also have octets. Multiple bonds are not needed, and the final structure is the one arrived at after Step 4.

Discussion Because we are using an algorithm to draw these structures, it is important that each step be carried out correctly because any errors may be propagated through the rest of the process. The first step is the most critical. If we lose track of the number of valence electrons we need to include, we will have little chance of arriving at the correct structure. Many students have missed exam questions by adding a lone pair of electrons they didn’t have rather than making multiple bonds.

Check Your Understanding Draw the Lewis structure for C2H6.


7. 5

Calcium phosphate is an important precursor for the formation of bioceramic coatings. Draw the Lewis structure of the phosphate ion, PO43–.

Strategy Follow the steps of the Lewis structure algorithm. Because phosphate is an anion, we’ll need to account for its charge in determining the number of valence electrons.

Solution Step 1: Count the total number of valence electrons. The phosphorus atom has five valence electrons, and each oxygen atom also has six. The ion as a whole carries a charge of 3–, which means it must contain three additional electrons. P

1×5= 5

O

4 × 6 = 24

+3 extra e− = 3 Total = 32

7.5


217

Step 2: Draw the skeletal structure of the molecule. Phosphorus is the least electronegative atom and appears first in the formula, so it should be the central atom. We’ll draw square brackets around the entire ion and indicate its charge. 3⫺

O O

P

O

O Step 3: Place single bonds between all connected atoms. 3⫺

O O

P

O

O Step 4: Place the remaining valence electrons not accounted for in Step 3 on individual atoms until the octet rule is satisfied. By drawing bonds to the oxygen atoms, we have accounted for 8 of the 32 electrons, and 24 remain to be assigned. Placing three lone pairs around each oxygen atom will account for all of the available electrons. 3⫺

O O

P

O

O Step 5: Create multiple bonds for any atoms that do not have a full octet of valence electrons. We have fulfilled the octet rule and have accounted for all of the available electrons. Multiple bonds are not required in this ion.

Discussion This is similar to the earlier examples, but the fact that we are working with an ion requires extra care in determining how many valence electrons to include.

Check Your Understanding Draw Lewis structures for the NH 4+ and ClO2– ions.

E XAM PL E PRO BLEM

7.6

Poly(vinyl alcohol) is used in several biomaterials applications, including surgical sutures. Draw the Lewis structure of vinyl alcohol, CH2CHOH, the monomer from which poly(vinyl alcohol) is made.

Strategy Follow the steps of the Lewis structure algorithm. Because we have a larger molecule, we will have multiple central atoms in this case.

Solution Step 1: Count the total number of valence electrons. Each carbon atom has four, each hydrogen atom has one, and the oxygen atom has six. So there are 18 valence electrons. Number of valence electrons = 2(4) + 4(1) + 1(6) = 18 Step 2: Draw the skeletal structure of the molecule. This is an organic compound, and our knowledge of organic chemistry suggests that the carbon atoms will be in the center of the molecule and will be linked by a chemical

218


bond. So we’ll place both carbon atoms in the middle, surrounded by the remaining atoms. We also saw in Chapter 2 that alcohols contain the !OH functional group, so we should expect vinyl alcohol to show that feature. And we know that hydrogen is always an outer atom, forming just one chemical bond. Combining those ideas, we should have three hydrogen atoms and one !OH group bound to the carbon atoms. The way the formula is written above tells us that we want two hydrogens bound to one carbon atom and then a hydrogen and the hydroxyl group bound to the other.

H

H

H

C

C

O

H

Step 3: Place single bonds between all connected atoms. To hold the molecule together, we need at least one bond between each pair of adjacent atoms.

H

H

H

C

C

O

H

Step 4: Place the remaining valence electrons not accounted for in Step 3 on individual atoms until the octet rule is satisfied. Making 6 bonds requires 12 electrons, so we still have 18 – 12 = 6 electrons available. We cannot assign lone pairs to the hydrogen atoms. Oxygen is more electronegative than carbon, so we start by assigning two lone pairs to the oxygen atom so that it has an octet. There are two remaining valence electrons, which are placed for the moment on one of the carbon atoms. (At this point the choice of which carbon gets those electrons is arbitrary.)

H

H

H

C

C

O

H

Step 5: Create multiple bonds for any atoms that do not have a full octet of valence electrons. The carbon atom that did not receive a lone pair in the last step does not have an octet, so we’ll need to convert a lone pair from an adjacent atom into a bonding pair, creating a double bond. But should this double bond be formed between the two carbon atoms or between the carbon atom and the oxygen atom? At least two factors point toward the carbon–carbon choice. First, we have already seen that carbon often forms multiple bonds. Second, placing the double bond between the two carbon atoms gives the molecule a more symmetrical structure, and such arrangements are usually preferable.

H

H

H

C

C

O

H

H

H

H

C

C

O

The C"C double bond also leads to a structure in which each carbon atom forms four bonds, consistent with the rules we learned in Chapter 2 for writing line structures.

H

Discussion This example shows that using our algorithm for obtaining a Lewis structure sometimes leaves us needing to make choices. As you grow more experienced in chemistry, decisions such as where to place a double bond become easier. Often you can look for clues, as we did in this case. Carbon routinely forms double (or even triple) bonds. Further, we know that the presence of the carbon-carbon double bond allows the addition reaction that would be needed to produce the polymer from the monomer. But what happens when the choice is less obvious? Such a situation is discussed in the next subsection.

Check Your Understanding Draw the Lewis structure for carbon dioxide.

7.5


219

Chemists are quite fond of Lewis structures, and the reason for this is clear. The structures are generally simple to draw; yet they can often provide powerful insight into the bonding within a molecule. On occasion, however, when the final step of our algorithm is carried out, the choice of where to put a double bond appears entirely arbitrary. The concept of resonance has been developed to help describe the bonding in such molecules within the framework of Lewis structures.

Resonance Let’s consider sulfur dioxide, one of the EPA’s criteria pollutants we discussed in Chapter 5. What should its Lewis structure look like? Both sulfur and oxygen have six valence electrons, so we have a total of 18 valence electrons. Sulfur occurs first in the formula and is less electronegative, so we place it at the center of the skeletal structure. O

S

O

We next place bonding pairs between the sulfur atom and each oxygen atom. O

S

O

We still have 14 valence electrons to add, starting with the oxygen atoms. When we place all 14 electrons in the structure as lone pairs, we have the following structure: O

S

O

Checking for octets, we find that both oxygen atoms have filled octets, but the sulfur has only six electrons. It needs two more, which can be obtained by forming a double bond. But on which side should we place this double bond? Unlike Example Problem 7.6, here we are confronted with two equivalent choices for the position of the double bond. We could draw either of the two structures. O

S

O

O

S

O

Each of these structures has the same fundamental weakness: it implies that the central sulfur atom somehow interacts more strongly with one oxygen atom than with the other. The explanation for this contradiction is that the structure of the molecule cannot be accurately represented by a single Lewis structure. Instead we say that the real structure corresponds to an average of all the valid Lewis structures, and we call this average a resonance hybrid of the contributing Lewis structures. It is very important to emphasize that the double-headed arrow does not mean that the structure alternates between the two possibilities; it indicates that the two structures shown for SO2 are resonance structures, contributing to a resonance hybrid. In resonance structures, the positions of all atoms are identical; only the positions of the electrons are different. The key point to emphasize here is that a resonance hybrid has a structure that is an average of the contributing resonance structures. To explain this statement better, let’s take another look at the two possible structures of SO2 shown above. Either resonance structure taken alone would suggest that an SO2 molecule contains one single bond and one double bond. If this were true, then one bond (the single bond) should be longer than the other (the double bond). But bond lengths are experimentally measurable, and the data clearly show that the two bonds are equal in length. Moreover, the bond lengths measured for SO2 are intermediate between those typically seen for an S!O single bond and an S"O double bond. These experimental results confirm the idea that neither resonance structure alone gives a good picture of the SO2 molecule and reinforce the view that the real structure is best thought of as an average of the possible resonance forms. When we can draw resonance structures, the actual structure is a hybrid, an average of the contributing structures, and NOT a mixture of them. The bonds can be thought of as “spread out” or delocalized over the molecule. 220


H H

H

C C

C

C H

H C C

H

H

H

H

C C

C

C

C C

H

H

H

Figure 7.11 ❚ The two resonance structures of benzene show there are two ways to distribute three double bonds among the six carbon atoms. Benzene does not behave chemically like a compound with three double bonds, however, because the true structure is an average between these two resonance structures.

Thus for SO2, we might think of having two “one and a half bonds” rather than a single bond and a double bond. Given its simplicity, it is not really surprising that the Lewis dot model can’t adequately describe every molecule we come across. The concept of resonance can be thought of as a patch that attempts to rationalize the experimental facts, while preserving most of the simplicity of the Lewis model. We have considered the structural implications of resonance, but there are also energetic consequences of resonance. This means that molecules with resonance structures behave differently from what one might expect on the basis of the individual resonance structures, which show a combination of single and double bonds. Benzene is the classic example of a substance that displays resonance, and in fact the concept of resonance originated from a need to explain the chemistry of benzene. Benzene, C6H6, is a volatile liquid obtained from coal tar. It was accepted in the latter half of the 19th century that benzene is a cyclic compound, whose carbon atoms form a six-membered ring. The only structures that could be drawn for it, however, had alternating single and double bonds (Figure 7.11). If these structures really represented the structure of benzene, the substance should have shown the addition reactions typical of compounds with C"C bonds. But benzene is actually much less reactive than alkenes. This means that the molecule is more stable than any one Lewis structure would predict. Over time, it became apparent from examining many compounds that this result is more general: when a molecule can be represented by multiple resonance structures, it is more stable (i.e., has lower energy and is less reactive) than would be predicted on the basis of any one of the contributing resonance structures. There are ways to understand this energy reduction based on quantum mechanics, but we will not present them here.

7.6

The source of the energy differences in resonance is the kinetic energy part of the quantum mechanical model. The details of this treatment are not necessary for an entry-level chemistry course.

Orbital Overlap and Chemical Bonding

Lewis structures provide us with a convenient way to gain some insight into a molecule’s bonding. But they don’t really tell us much about how chemical bonds might be formed, other than through the general notion of electron sharing. Scientific curiosity leads us to want to know more about covalent bonds, as does a sense that understanding bonding will pay off in practical applications. Because we recognize shared electrons as the “glue” that holds covalent bonds together, we should expect any deeper explanation to focus on electrons. In Chapter 6, we learned that the best current model of the atom views electrons as delocalized waves, described by orbitals. If electrons behave as waves, then to explain how electrons can form bonds between atoms, we must consider how waves interact with one another. To explain how two waves!or two orbitals!might interact with one another, we will need to invoke the concept of interference, which provides a clear contrast between waves and particles. When two particles occupy the same space, they collide, but when two waves occupy the same space, they interfere. As you may have learned in a high school physics or physical science class, this interference may be either constructive or destructive. In constructive interference, two interacting waves reinforce one another to form a larger wave. In destructive interference, the two interacting 7.6 Orbital Overlap and Chemical Bonding

221

⫹ (a)

⫹ (b)

Figure 7.12 ❚ Simple sine waves illustrate the concept of interference between waves. In the upper panel (a), the two identical sine waves on the left are totally in phase with one another, meaning that their peaks and valleys line up. Adding these two waves will produce constructive interference, and the resulting wave will have twice the amplitude of the original waves. In the lower panel (b), the two waves are out of phase, so that the peaks of one correspond to the valleys of the other. Combining these waves will result in destructive interference, where the two waves effectively cancel one another and the result is a flat line.

The name valence bond model arises from the fact that only valence electrons are considered.

waves offset or “cancel” one another to produce a smaller wave or no wave at all. Both constructive and destructive interference are illustrated in Figure 7.12. In the quantum mechanical model, electrons are simply one type of wave, and we can picture the formation of chemical bonds as an example of constructive interference between electron waves. But for electron waves from different atoms to interact they must first occupy the same region of space. Another way of saying this is that the atoms must be positioned so that the valence orbitals from one atom must overlap those of the other atom if a bond is to form. This idea forms the basis for the valence bond model of chemical bonding, in which all bonds are seen as the result of overlap between atomic orbitals. Let’s examine this idea of orbital overlap by looking at the simplest possible molecule, H2. Hydrogen atoms have a single valence electron in a 1s orbital. To form a diatomic hydrogen molecule, the 1s orbitals from two different hydrogen atoms must overlap. Figure 7.13 shows what this overlap might look like and how interference between the wave functions of the two electrons produces a bond between the atoms. Note that because 1s orbitals are spherical, there is no preferred direction of approach. But when p orbitals are involved, their more complex shape will lead to different geometries for

H—H

Figure 7.13 ❚ The covalent bond in molecular hydrogen can be described in terms of the overlap between the 1s orbitals from the two H atoms. The upper diagram in the figure shows this overlap by plotting the wave functions for the 1s orbitals. Constructive interference between these wave functions in the region between the nuclei results in bond formation. The lower diagram uses shading to represent the same buildup of electron density.

222


Figure 7.14 ❚ When two p orbitals approach one another and overlap end to end, as shown here, they form a sigma bond. (This is sometimes denoted by the lowercase Greek letter sigma: s.)

orbital interaction. These different geometries can be used to distinguish between types of chemical bonds. Using Lewis dot structures, we can show easily that N2, for example, must have a triple bond between the two nitrogen atoms. Although this large buildup of negative charge density between the nuclei allows very strong attraction between them, it also raises a question. How do six electrons occupy the same region of space without strongly repelling each other? We can address this question by considering the shape of the orbitals involved and their arrangement as overlap occurs. The valence electron configuration for a nitrogen atom is 2s22p3. Because the 2p orbitals are all singly occupied, overlap of a 2p orbital from one atom with a corresponding 2p orbital from the other atom should enable bond formation. But because of their shapes, p orbitals can overlap in different orientations. Recall that the three 2p orbitals from any atom can be described as lying along the x, y, and z axes. (See Figure 6.13, p. 179.) Let’s assume that two nitrogen atoms approach one another along the z axis. As seen in Figure 7.14, the 2pz orbitals can overlap end-to-end. This leads to a buildup of electron density along a line drawn between the two nuclei, just as we saw for H2. This type of constructive interference produces what are called sigma bonds. The accumulation of electron density in a sigma bond sits along the line connecting the two nuclei. This type of overlap can also occur in bonds in which an s orbital overlaps a p orbital or when two s orbitals overlap. Now let’s consider the 2px and 2py orbitals of the nitrogen atoms. As the nuclei draw closer together, these orbitals can overlap in a side-to-side orientation, as shown in Figure 7.15. Once again the overlap leads to constructive interference. This allows the buildup of electron density between nuclei that is crucial to chemical bonding. When this sideways overlap occurs, however, the buildup of electron density will not be localized along the line between the nuclei. As shown in Figure 7.15, the electron density is localized above and below or in front and in back of this line. Such bonds are called pi bonds. In the N2 molecule, then, we can envision one pi bond formed by the overlap of the 2px orbitals from each atom and another pi bond formed by the overlap of the 2py orbitals. The triple bond thus comprises one sigma bond and two pi bonds. The idea that bonding involves the overlap and interaction of orbitals from different atoms is a fundamental building block for a more detailed understanding of the chemical bond. For this reason, the valence bond model continues to play an

N

N

Figure 7.15 ❚ When two p orbitals approach one another and overlap side to side, as shown here, they form a pi bond. (This is sometimes denoted by the lowercase Greek letter pi: p.) In the example shown, orbital overlap will produce a buildup of electron density above and below the line defined by the two nuclei.

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223

important role in the way scientists envision bonding. We do, however, encounter serious difficulties in describing bonding through the overlap of atomic orbitals. Because p orbitals form 90º angles with one another, we might predict that many molecules, including H2O and OF2, would show 90º bond angles. But such 90º angles are quite rare. Furthermore, looking at the electron configuration of carbon, we might expect compounds such as CH2 or CF2 to be stable and commonplace. But instead we find CH4 and CF4, and the simple overlap of atomic orbitals cannot help us to rationalize this. So we must extend our valence bond model to allow more extensive interactions between orbitals.

7.7

Hybrid Orbitals

Our notion of orbitals arose from the quantum mechanical treatment of individual atoms. Despite the obvious conceptual challenges that quantum mechanics poses, we accept the existence of s, p, d, and f orbitals because, among other things, doing so provides such a reasonable way to understand the shape of the periodic table. When atoms combine to form molecules, however, the observations noted above suggest that often those same orbitals cannot explain the observed facts. We just noted that electrons in p orbitals would be expected to lead to an H!O!H bond angle of 90º in water, for example, and this does not agree with the observed geometry of the molecule. And yet the relative simplicity of the idea that bonds result from orbital overlap remains alluring. This leads to the introduction of the concept of orbital hybridization, which is intended to reconcile the notion of orbital overlap with a wide range of observations on molecular shapes and structures. At its heart, the idea of hybridization acknowledges the fact that when two or more atoms approach one another closely enough to form chemical bonds, the interactions can be strong enough to reshape the orbitals on those atoms. The simplest stable binary compound formed by carbon and hydrogen is methane, CH4, for which the Lewis structure is shown below along with a molecular model.

H H

C

H

H

Orbital theory began as a model. As long as adjustments, such as the invention of hybrid orbitals, do not change the premise of the model, such changes can be made—and are worth making if they provide new insight.

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Bond lengths and strengths can be measured experimentally, and such measurements clearly show that all four C!H bonds in CH4 are identical. Furthermore, molecular shapes can also be measured by observing the rotation of molecules, and such measurements show that the H!C!H bond angles in CH4 are all 109.5º, as shown in the molecular model above. These data suggest that if we wanted to describe the bonds in CH4 in terms of orbital overlap, we would need four identical orbitals on the carbon atom and that these orbitals should be arranged at angles of 109.5º from one another. When we introduced the idea of orbitals, we noted that they were mathematical in origin. In mathematics, we can manipulate functions and transform them according to accepted rules. The wave functions for atomic orbitals can also be manipulated in the same ways. We need not be concerned with all the details of such manipulations, but one rule will be important; the number of orbitals we have must not change due to the manipulation. If we have two orbitals to begin with, we must still have two orbitals when we finish. Among other things, this ensures that we can still accommodate all the electrons needed. If we take linear combinations of atomic orbitals, we can devise new orbitals that can be used as the basis for a model of bonding. The resulting orbitals are called hybrid orbitals because they are mixtures of two or more atomic orbitals.


Energy, E

Figure 7.16 ❚ The box diagram 2p 2p 2p Hybridization

sp3

depicts the energetics of sp3 hybridization. The sp3 hybrid orbitals have lower energy than the original p orbitals but higher energy than the original s orbitals.

2s

One way to envision this process is to use boxes to represent orbitals, as we did in Chapter 6 when we introduced electron configurations. Figure 7.16 shows this for the hybridization of a carbon atom. The s orbital and the three p orbitals are mixed to form four new hybrid orbitals, all having the same energy. The name given to these orbitals is sp3, where the superscript indicates that three p orbitals are included. Placing electrons in these four equivalent sp3 orbitals means that each will be singly occupied. Thus these orbitals readily explain the ability of carbon to form four identical bonds. If we were to examine the details of the mathematics, we would also discover that each of these sp3 orbitals is strongly oriented in a particular direction at angles of 109.5º between orbitals. These hybrid orbitals thus make it easy to explain the observations for CH4. The overlap of an sp3 orbital from carbon and an s orbital from hydrogen forms each of the four C!H bonds. The concepts behind the hybrid orbital model are very flexible with regard to the number of orbitals that can be combined, as shown in Table 7.2. When only one Table

❚

7.2

The five common orbital hybridization schemes are shown. The names for these hybrid orbitals are derived from the type and number of orbitals that combine to form them. Thus sp2 hybrids result from combining an s orbital with a pair of p orbitals. The orbital geometries shown in the left-hand column give rise to the common molecular shapes described in the next section.

Orbitals Combined

Hybridization

Orbital Geometry

s, p

sp

s, p, p

sp2

A

s, p, p, p

sp3

A

s, p, p, p, d

sp3d

s, p, p, p, d, d

sp3d 2

Hybridization with five or six orbitals also allows us to explain the fact that some elements form compounds with expanded octets—more than eight electrons around a single atom. In the figures shown here, each lobe represents a different hybrid orbital. This is an important contrast to the unhybridized p orbitals, where each orbital consists of two lobes of electron density.

A

A

A

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Hybrid Orbitals

225

p orbital is included in the scheme, we form sp hybridized atoms, and when two p orbitals are included, the hybrid orbitals are sp2. For atoms with valence electrons in the n = 3 shell or higher, hybridization can also include d orbitals. Thus sp3d hybridization yields five orbitals, and when six orbitals are included, we have sp3d2 hybrids. All of these schemes can be used to help explain bonding in covalent molecules.

7.8

If you don’t have access to molecular models, you can build your own by using small balloons and holding them together with a clip of some sort. The balloons must stay out of each other’s way physically, analogous to the electronic repulsions of electrons.

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Shapes of Molecules

Although Lewis structures allow us to account for the number of bonds in a molecule and show us the bonding pattern (which atoms are joined together and by how many bonds), they give us an incomplete, two-dimensional picture of the molecular structure. If we extend our analysis of the molecular structure to three dimensions, we need to examine the order in which the atoms are joined to one another and also how they are arranged in space—in other words, the molecular shape. In addition to being of theoretical interest, the shapes of molecules are important because they affect molecular properties, including reactivity. Being able to predict the shapes of molecules is crucial to understanding the properties and reactivity of a polymer like PMMA, the polymeric proteins and nucleic acids that are the operating systems in living organisms, the catalysts used in industrial processes, and the molecules synthesized for medicines and consumer goods. One thing that sets experienced chemists apart from beginning students is the ability to read a molecular formula and visualize what the molecule would look like on the microscopic scale. To arrive at an understanding of the shapes of molecules, we need to analyze Lewis structures in more detail. A very useful technique for doing this relies on the valence shell electron pair repulsion (VSEPR) theory, which states that molecules assume a shape that allows them to minimize the repulsions between electron pairs in the valence shell of the central atom. Lewis structures emphasize the importance of pairs of electrons around atoms, either as bonding pairs or as lone pairs. If we consider that these pairs occupy localized regions of space, we would expect that the various regions would be as far away from each other as possible. After all, these regions are centers of negative charge, and like charges repel one another. Couple this electron pair repulsion with the fact that Lewis structures include only electrons in the valence shell, and the origin of the name VSEPR is apparent. The theory provides a very simplistic view of the electrons in molecules, but it leads to surprisingly accurate predictions of molecular geometries. VSEPR is an application of concepts from geometry. In particular, we want to answer this question: If the bond lengths are fixed, what geometric shapes allow the regions of negative charge represented by various electron pairs to be as far away from one another as possible? The answer depends on the number of electron pairs that surround the atom. Table 7.3 shows the basic geometry for distributing two through six electron pairs to minimize repulsions. Examining this table, you should realize that most of the shapes involved are fairly simple. It should be intuitively clear that when only two regions of electron density are present, for example, a linear arrangement would maximize their separation. Of all the shapes shown, the tetrahedral and trigonal bipyramidal arrangements are probably the least familiar. You should be sure that you have a real understanding of what these different geometries look like. (Models!either physical or software-based!can be very helpful.) Most of the bonding arrangements we will encounter can be analyzed in terms of the five basic shapes shown here. We can predict molecular geometries systematically. Because our approach centers on the arrangement of electron pairs around the central atom, we will always start by drawing a Lewis structure. Once we’ve done that, it is simple to count the number of bonding and nonbonding electron pairs around the central atom. We’ll start with the most straightforward examples, in which all the electrons around the central atom


Table

❚

7.3

Each of the geometrical arrangements shown in the table minimizes the electron pair repulsions for the indicated number of electron pairs. To visualize the shapes of molecules, it is essential that you have a sound mental picture of each of these geometries.

Number of Electron Pairs

Geometric Name

Bond Angles

Diagram 180°

2

Linear

180º 2: Linear

120°

3

Trigonal planar

120º

3: Trigonal planar

109.5°

4

Tetrahedral

109.5º

4: Tetrahedral

90°

5

Trigonal bipyramidal

120°

120º, 90º

5: Trigonal bipyramidal

90°

6

Octahedral

90°

90º, 180º

6: Octahedral

are in single bonds. Then we’ll find out how to determine the molecular geometry when there are also nonbonding pairs or multiple bonds involved. If the central atom is surrounded only by single bonds, the process of determining the molecular geometry is fairly simple. We count the number of single bonds emanating from the central atom and use the information in Table 7.3 to correlate that number of electron pairs with the geometry predicted by VSEPR, and this describes the shape of the molecule.

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E XAM PL E PRO BLEM

7.7

Determine the shape of each of the following species: (a) PO3– , (b) PCl5 4

Strategy Draw the Lewis structures. Each of these molecules has only single bonds around its central atom. So we can count the number of bonding pairs around the central atom and assign the geometry by consulting Table 7.3 as needed.

Solution (a) The Lewis structure of PO3– , which we drew in Example Problem 7.5, is 4 3⫺

O O

P

O

O The central carbon atom has a complete octet of electrons, contributed by four bonding electron pairs. The four electron pairs have minimal interaction when they point toward the corners of a tetrahedron. The ion is tetrahedral (Table 7.3). (b) The Lewis structure of PCl5 is somewhat unusual because the central phosphorus atom has more than an octet around it. There are a number of molecules with this property and they serve as useful examples for structures in VSEPR. Cl Cl Cl

P

Cl

Cl The central phosphorus atom has an expanded octet with 10 electrons arranged in five bonding pairs. The five electron pairs occupy the vertices of a trigonal bipyramid (Table 7.3). Note that, unlike the oxygen atoms in the tetrahedral ion we looked at in (a), all of the chlorine atoms in the trigonal bipyramid are not in equivalent positions. There is a spacing of 120º between the chlorine atoms arranged in the horizontal triangle (the equatorial positions), whereas the two atoms above and below the plane of the triangle (axial positions) form an angle of 90º with the plane of the triangle.

Discussion These examples can be difficult to see using only the text. Model kits or software models can help you visualize the three-dimensional nature of the shapes described here.

Check Your Understanding Predict the shape of the SF6 molecule. Whenever we talk about the geometry or shape of a molecule, we refer to the positions of its atoms. This seems obvious in cases like those in Example Problem 7.7, where there are no lone pairs on the central atom. But the same idea holds, even if there are lone pairs involved. The positions of the atoms define the shape of a molecule. But lone pairs also contribute to the overall repulsion between electron pairs, and so they do influence the geometry of molecules. When the central atom is surrounded by both bonding and nonbonding electron pairs, we can choose a base geometry from the list in Table 7.3 by counting the total number of regions of high electron density!bonding pairs plus lone pairs. To determine the actual molecular geometry, we visualize what that shape would look like if the positions to be occupied by lone pairs were empty. In some cases, more than one shape may

228


seem possible. When that happens, the guiding premise of VSEPR is that the correct shape will be the one that minimizes the overall repulsion between electron pairs. In deciding which structure that is, the order of repulsive interactions is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. This order can be explained in terms of the fact that lone pairs are attracted only to one nucleus. This means that a lone pair is less tightly localized, so it occupies a larger space. Examples of the various possible geometries are shown in Table 7.4. Determining the geometry of a molecule with both bonding and nonbonding electrons is best illustrated by some examples.

Table

❚

7.4

The molecular shapes resulting from various combinations of the total number of electron pairs around the central atom and the number of lone pairs


Number of Lone Pairs

Shape

3

0

Trigonal planar

3

1

Bent (120º)

4

0

Tetrahedral

4

1

Trigonal pyramidal

4

2

Bent (109.5º)

5

0

Trigonal bipyramidal

Ball and Stick Model

Note that as long as you have a good grasp of the basic geometries (Table 7.3), it is not necessary to memorize this extensive listing. Instead you should be able to start from the basic geometry for the number of electron pairs present and visualize the shape that would result from “emptying” a number of positions corresponding to the lone pairs.

(continues)

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Shapes of Molecules

229

Table

❚

7.4 (continued)

The molecular shapes resulting from various combinations of the total number of electron pairs around the central atom and the number of lone pairs


230

Number of Lone Pairs

Shape

5

1

Seesaw

5

2

T-shape

5

3

Linear

6

0

Octahedral

6

1

Square pyramidal

6

2

Square planar

6

3

T-shape

6

4

Linear


Ball and Stick Model


7. 8

Determine the shape of the following molecules using VSEPR theory: (a) SF4, (b) BrF5

Strategy As always, we start by drawing the Lewis structures. Then count the number of electron pairs around the central atom and determine the spatial arrangement of electrons pairs, consulting Table 7.3 as necessary. Place the lone pairs in positions where the electron repulsions are minimized and describe the resulting geometric arrangement of the atoms.

Solution (a) The Lewis structure of SF4 is as shown. F

F S

F

F

There are five pairs of electrons around sulfur, so the shape with minimum interaction is a trigonal bipyramid. Unlike the previous example, though, here the central sulfur atom has one lone pair. The five positions of a trigonal bipyramid are not equivalent, so we need to determine the best location for the lone pair. Electron pairs in the planar triangle (equatorial positions) have two near neighbors at an angle of 120º and two at 90º, whereas the axial positions have three near neighbors, each at an angle of 90º. Thus there is “more space” for electron pairs in the equatorial positions. Because lone pair–bond pair repulsion is larger than bond pair–bond pair repulsion, this means that the lone pair should occupy an equatorial position. This reduces the overall repulsive interactions. (Only the two closest bonding pairs are 90º away, whereas if the lone pair was in an axial position, there would be three bonding pairs 90º away.)

The resulting shape is usually described as a seesaw. If you don’t see why, try looking at a model. (b) The Lewis structure of BrF5 is as shown. F

F

F

Br F

F

There are six electron pairs around the central bromine atom, so the distribution of electrons is based on an octahedron. A lone pair occupies one of the six positions, so we must consider where to place it. But because an octahedron is completely symmetrical and all of its vertices are equivalent, the lone pair can occupy any corner. The resulting shape of the molecule is a square pyramid.

Discussion For many students, the confusing part of this type of problem lies in the appearance that a lone pair of electrons is considered for part of the problem but not for all of it. There are two things to remember: (1) The shape of a molecule is determined

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231

by the positions of its atoms. This is a result of the ways we measure shapes. Because nuclei are so much more massive than electrons, our experiments detect them more readily, and thus we consider that the shape is defined by the positions of the nuclei. (2) Whether or not it is involved in bonding, every electron pair in the Lewis structure constitutes a region in space where there is a relatively high concentration of negative charge. The key concept of VSEPR theory is that these negatively charged regions will repel each other and therefore must be arranged as far apart as possible. So the electrons dictate the shape of a molecule, but they are not directly measured. That’s why we have a process where we appear to consider them at some stages but not at others.

Check Your Understanding What are the shapes of the following molecules? (a) ClF3, (b) XeF2

In determining the geometry of molecules by VSEPR, multiple bonds are treated as a single region of electron density. This is intuitively reasonable because the position of greatest electron density for a multiple bond lies in the region between the two bonding atoms, just as for a single bond. To put this in another way, it wouldn’t make sense to say that the electron pairs making up a double or triple bond should point in totally different directions from one another. Thus for purposes of determining molecular shapes, any double or triple bonds behave as if they were a single electron pair.

E XAM PL E PRO BLEM

7.9

Use VSEPR theory to determine the shape of the NOF molecule.

Strategy Once again, we start by drawing the Lewis structures. Then count the number of regions of electron density around the central atom, remembering to count any double or triple bonds as a single region. Determine the spatial arrangement of electrons pairs, consulting Tables 7.2 as needed. Place any lone pairs in positions where the electron repulsions are minimized and describe the resulting geometric arrangement of the atoms.

Solution The choice of central atom is perhaps less obvious here than in some cases we have looked at. But nitrogen is the least electronegative atom in this molecule, so it will be central. Starting from an F!N!O skeleton, the Lewis structure is as shown below. O

N

F

There are three regions of electron density around the nitrogen atom, so the shape with minimum interaction is trigonal planar. The central nitrogen atom has one lone pair. In the trigonal planar geometry, all three positions are equivalent, so it does not matter which one the lone pair occupies. The molecule will be bent, with an O!N!F bond angle of approximately 120º. Because lone pairs provide stronger repulsion, the actual bond angle should be slightly less than 120º.

Discussion Much like lone pairs, double bonds can cause confusion for many students. When counting the number of valence electrons to draw the Lewis structure, we must be sure to count the double bond as four electrons. (And a triple bond, if one were present, would count as six electrons.) But in determining the shape, we treat the double (or triple) bond as a single region of electron density. It may help if you realize that both of the bonds in the double bond must “point” in the same direction because they connect the same pair of atoms.

232


Check Your Understanding What are the shapes of the following molecules? (a) COF2, (b) COS Applying the same principles as for smaller molecules, VSEPR theory can also be used to predict the geometry of larger molecules. In predicting the structure of a molecule with more than one “central” atom, we look at each atom within the backbone of the molecule and predict the bonding about that atom, using the same procedures that we have just outlined. For example, let’s see what happens when we apply VSEPR theory to the structure of vinyl alcohol, the monomer of polyvinyl alcohol, H2C=CHOH, which has important biomaterials applications such as in surgical sutures. The Lewis structure we drew in Example Problem 7.6 is a good place to start.

H

H

H

C

C

O

The order in which we look at different central atoms is not important.

H

Each of the two carbon atoms is bound to three other atoms and has no lone pairs, implying trigonal planar geometry. Next, we need to ask whether or not those two trigonal planar ends of the molecule are coplanar. Here, it helps to invoke hybridization and think about the orbitals involved in forming the double bond. Each can be described as using sp2 hybrid orbitals. For each carbon atom, two of those hybrid orbitals will overlap with orbitals from the hydrogen or oxygen atoms, and the third overlaps an sp2 orbital from the other carbon. This gives us a single sigma bond between the carbon atoms so far. But each carbon also still has one remaining unhybridized p orbital, and that p orbital is perpendicular to the plane of the sp2 hybrids, as shown in Figure 7.17. If these p orbitals are aligned parallel to one another, they will be able to overlap side to side, forming a pi bond. This explains the existence of the double bond and also tells us that the carbon atoms and the atoms bound to them must all lie in the same plane. If one end of the molecule were to rotate with respect to the other, the p orbitals would no longer be in the right orientation to form the pi bond. p orbitals; π bond

H

H

C

C

sp3 hybrid orbitals; σ bonds

H H s orbital; σ bond

O

Figure 7.17 ❚ The orbitals involved in the bonding in H2C"CHOH are shown. The s-bonding

framework results from the overlap of sp2 hybrid orbitals on the carbon atoms with (a) each other, (b) the s orbitals on hydrogen atoms, and (c) an sp3 orbital on the oxygen. The p bond is due to overlap of the remaining unhybridized p orbitals of the carbon atoms. Thus the electron density of the p bond would be above and below the plane of the paper. The shape at each carbon atom is trigonal planar, and the shape at the oxygen atom is bent. In the view shown here, all seven atoms lie in the plane of the paper, and the lone pair electrons on the oxygen are above and below the paper. But the H-atom in the OH group is actually free to rotate about the C!O bond axis.

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233

INSIGHT INTO

7.9

Nanotechnology will be discussed further in Chapter 8.

Molecular Scale Engineering for Drug Delivery

Throughout this chapter, we have seen a number of ways in which aspects of chemical bonding are important factors in the design of materials for biomedical engineering applications. It is increasingly possible to consider doing engineering design at the molecular level, and one active area for this work lies in more targeted delivery of drugs for therapeutic benefits. Many pharmaceuticals, particularly those used in the treatment of cancer through chemotherapy, have significant side effects. One reason why side effects occur is because the drug molecules affect healthy cells as well as diseased ones. Thus a strategy that uses new biomaterials to deliver drugs exclusively to the desired locations (in tumor cells, for example) would represent an important advance in biomedical engineering. This must be accomplished within the constraints of biocompatibility that have already been emphasized in this chapter. A promising method for this type of drug delivery system uses a material called mesoporous silica nanoparticles (MSN). As we saw earlier, silica is composed of networks of SiO4 units, and those SiO4 units can form a honeycomb structure as shown in Figure 7.18. An MSN is simply a very small particle with this honeycomb arrangement. Because of this structure, these particles have enormous surface to volume ratios: one gram of the material has roughly the same surface area as a football field. Moreover, the pores in a MSN can be used to store drug molecules. Once loaded with the desired therapeutic agents, the pore can be capped with another molecule, and the whole nanoparticle is delivered to the target. Earlier we noted that amorphous silica particles can destroy red blood cells and so are not biocompatible. In light of this knowledge, it may seem surprising that these MSNs could be biocompatible. The difference arises from the honeycomb structure of the MSNs. Although the overall surface area of the MSN is very large, the majority of that surface area represents the interior walls of the honeycomb pores. Relatively little of the surface area is actually on the outside of the particle itself. And since the pores are smaller than red blood cells, those SiO4 groups lining the inner walls of the pores cannot damage blood cells. The development of nanoparticles like these MSNs for drug delivery represents an engineering design project at the molecular scale. Perhaps the most critical element

Pacific Northwest National Laboratory

Figure 7.18 ❚ This drawing shows the honeycomb structure found in mesoporous silica. The insets show how an individual pore might be loaded with small molecules in a drug delivery application.

234


in this design is how the cap molecules at the end of the pores are to be held in place. The materials chosen for this task must be capable of staying in place until an agent in a target cell somehow breaks the bond. To undertake this type of work, biologists, engineers, and chemists work together to tune the type and strength of bonding used, showing once again how the collaboration between experts in different fields moves modern technology forward.

FO C U S ON


Question An unidentified solid is dissolved in water to produce a clear, colorless solution, and this solution conducts electricity. A second solution is added, and a precipitate forms. Once the precipitate has settled, the liquid above it does not conduct electricity perceptibly. What does this experiment tell you about the bonding in the initial solid and in the precipitate, and how does it tell you that?

Strategy This question is another example of conceptual problem solving. We have seen three types of chemical bonding in this chapter, and from the observable behavior of this experiment, we need to infer which one is important in the unknown solid. The other piece of information we need to include involves what allows a solution to conduct electricity, as discussed in Chapter 3.

Answer For any substance to conduct electricity, it must allow a charge to move. The charged particles in solution are usually ions, so this flow of charge arises from the movement of ions. This strongly suggests that the original solid was ionically bonded. How does this explain the observations? When the ionic solid is dissolved, its constituent ions dissociate from one another, allowing the solution to conduct electricity. The addition of the second solution forms a precipitate. This tells us a couple of things. First, because it was formed from two solutions, the precipitate must be an insoluble ionic compound. Moreover, one of the constituent ions in the precipitate must have come from each solution. If the anion came from the first solution, the cation must have come from the second. We can also note that because the final solution no longer conducts electricity, it must not contain significant concentrations of any ions. This points toward a few possibilities. Each of the individual solutions must have contained both anions and cations to be electrically neutral. When the two solutions were mixed, these ions must have recombined to form two products: the cation from the first solution combining with the anion from the second and vice versa. So we have two products to account for. One possibility is that two different precipitates may have formed, removing all of the ions. Alternatively, the second product could be water, formed from the neutralization reaction between an acid and a base.

S U M M ARY Building on our understanding of atomic structure, we have now established a better picture of how atoms “stick together” in chemical bonds. The forces involved are those between charged particles, as is most easily seen in ionic bonding. When a metallic atom forms a cation and a nonmetal forms an anion, it is not at all surprising that these oppositely charged particles attract each other and produce what is referred to as an ionic bond. We can go further in our understanding of ionic bonds by viewing bond formation as a series of steps. Forming a cation requires an energy input equal to the ionization energy, and the formation of an anion generally releases energy, as indicated by the negative electron affinity. Finally, the arrangement of the oppositely charged particles into a lattice results in significantly stronger

attractive forces than repulsive ones, so metals and nonmetals routinely form ionic compounds. When bonds form between nonmetal atoms, electrons are shared in pairs rather than wholly transferred to form ions. This sharing of electron pairs is called covalent bonding. Different atoms possess varying capacities to attract electrons within a covalent bond, as measured by the electronegativity. So the distribution of shared electrons is not always symmetrical and leads to the idea of polar bonds. There is a continuum of behavior from even sharing (as between two identical atoms) in covalent bonds, to uneven sharing to form polar covalent bonds, and to complete transfer of electrons from one atom to another to form ionic bonds.

Summary

235

The propensity of atoms to share electrons can be summarized visually in a Lewis structure. By following a fairly simple set of rules, we can draw Lewis structures for many ions or molecules. We can then use those Lewis structures to predict the nature of the bonding involved. Most importantly, Lewis structures help us understand whether one or more pairs of electrons need be shared!leading to predictions of single versus double or triple bonds for molecules. Beyond predicting what types of bonds are present in a molecule, however, the Lewis structure tells us fairly little about the details of a chemical bond. To understand just how electrons are shared, we must realize that electrons behave as waves and that when waves overlap, they interfere with each

other. When the waves “buildup” they are said to interfere constructively and the chemical implication is the formation of a chemical bond. Overlap can be achieved in more than one way, so that we can distinguish types of bonds as either sigma or pi bonds. Finally, in addition to simply representing a pair of shared electrons, a chemical bond has structural implications as well. Because electrons are negatively charged, when there are several distinct bonds, they will tend to be physically separated from each other. This idea is the basis for a method to predict the geometry of molecules called the Valence Shell Electron Pair Repulsion (VSEPR) theory. Using this theory, the general shape of molecules and ions can be predicted.

KEY TERMS biocompatibility (7.1)

hybridization (7.7)

pi bond (7.6)

bond energy (7.3)

ionic bond (7.2)

polar bond (7.4)

bond length (7.3)

lattice energy (7.2)

resonance (7.5)

bonding pair (7.3)

Lewis dot symbol (7.3)

sigma bond (7.6)

covalent bond (7.3)

Lewis structure (7.3)

triple bond (7.3)

diploe (7.4)

lone pair (7.3)

valence bond model (7.6)

double bond (7.3)

molecular shape (7.8)

VSEPR theory (7.8)

electronegativity (7.4)

octet rule (7.3)

hybrid orbitals (7.7)

orbital overlap (7.6)



7.10

■ Arrange the members of each of the following sets of cations in order of increasing ionic radii: (a) K+, Ca2+, Ga3+, (b) Ca2+, Be2+, Ba2+, Mg2+, (c) Al3+, Sr2+, Rb+, K+, (d) K+, Ca2+, Rb+

7.11

■ Arrange the following sets of anions in order of increasing ionic radii: (a) Cl–, P3–, S2–, (b) S2–, O2–, Se2–, (c) Br–, N3–, S2–, (d) Br–, Cl–, I–

INSIGHT INTO Materials for Biomedical Engineering 7.1 Define the term biocompatibility. 7.2 List some properties associated with biomaterials used for joint replacements. 7.3 Describe how PMMA functions as bone cement. How does chemistry play a role in this process? 7.4 The formulas of sodium phosphate and sodium orthosilicate are Na3PO4 and Na4SiO4 respectively. What are the chemical formulas of calcium phosphate and calcium silicate? 7.5 Why do biomedical engineers sometimes need to use coatings on materials they use? When coatings are generated, what are the materials used as reactants called? 7.6 Use the concept of polarity of water and the basic composition of the body to explain why the polarity of biomaterials is important. 2+

7.7 Why is the Na

ion not found in nature?

7.8 Why do nonmetals tend to form anions rather than cations?

236

7.13 In a lattice, a positive ion is often surrounded by eight negative ions. We might reason, therefore, that the lattice energy should be related to eight times the potential of interaction between these oppositely charged particles. Why is this reasoning too simple? 7.14 Use the concept of lattice energy to rationalize why sodium fluoride dissolves in water, whereas calcium fluoride does not. Extending this reasoning, would you expect magnesium fluoride to be soluble? The Covalent Bond

The Ionic Bond

7.9

7.12 Which pair will form a compound with the larger lattice energy: Na and F or Mg and F? Why?

Select the smaller member of each pair: (a) N and N3–, (b) Ba and Ba2+, (c) Se and Se2–, (d) Co2+ and Co3+ ■


7.15 Describe the difference between a covalent bond and an ionic bond. 7.16 Considering the potential energy of interacting particles, how does covalent bonding lower the energy of a system?

■


7.17 Sketch a graph of the potential energy of two atoms as a function of the distance between them. On your graph, indicate how bond energy and bond distance are defined. 7.18 When a covalent bond forms, is energy absorbed or released? Explain how your answer is related to the graph you sketched in the previous problem.

7.34

on the positions in the periodic table of the following pairs of elements, predict whether bonding between the two would be primarily ionic or covalent. Justify your answers. (a) Ca and Cl, (b) P and O, (c) Br and I, (d) Na and I, (e) Si and Br, (f ) Ba and F.

7.35

■ In each group of three bonds, which bond is likely to be the most polar? Which will be the least polar? (a) C!H, O!H, S!H, (b) C!Cl, Cl!Cl, H!Cl, (c) F!F, O!F, C!F, (d) N!H, N!O, N!Cl

7.19 Coulombic forces are often used to explain ionic bonding. Are coulombic forces involved in covalent bonding as well? Explain. 7.20 In terms of the strengths of the covalent bonds involved, why are combustion reactions exothermic? 7.21 If the formation of chemical bonds always releases energy, why don’t all elements form dozens of bonds to each atom? 7.22 Draw the Lewis dot symbol for each of the following atoms. (a) boron, (b) fluorine, (c) selenium, (d) indium 7.23 Theoretical models for the structure of atomic nuclei predict the existence of superheavy elements that have not yet been discovered and suggest that such elements might be fairly stable if they could be produced. So researchers are currently trying to synthesize these superheavy elements to test these theories. Suppose that element 117 has been synthesized and given the atomic symbol Lt. What would the Lewis dot symbol be for this new element? 7.24 Use Lewis dot symbols to explain why chlorine bonds with only one hydrogen atom.

7.36 Considering the trends in electronegativity shown in Figure 7.8, explain why alloys that form between two or more transition metals are not ionic substances. 7.37 Can fluorine atoms ever carry a partial positive charge in a molecule? Why or why not? 7.38

How many electrons are shared between two atoms in (a) a single covalent bond, (b) a double covalent bond, and (c) a triple covalent bond?

■

7.27 How does the bond energy of a double bond compare to that of two single bonds between the same elements? How does this relationship explain the types of reactions that compounds with double bonds undergo? Electronegativity and Bond Polarity

Which of the following contains both ionic and covalent bonds in the same compound? (a) BaCO 3, (b) MgCl 2, (c) BaO, (d) H2S, (e) SO2– 4 ■

7.39 Using only halogen atoms, what would be the most polar diatomic molecule that could be formed? Explain your reasoning. 7.40 Suppose that you wanted to make a diatomic molecule using two different halogen atoms. What combination of halogens would be the least polar? Keeping Track of Bonding: Lewis Structures 7.41

■ Draw the Lewis structure for each of the following molecules: (a) CO, (b) H2S, (c) SF6, (d) NCl3

7.42

■ Draw a Lewis structure for each of the following molecules or ions: (a) CS2, (b) BF4–, (c) HNO2 (where the bonding is in the order HONO), (d) OSCl2 (where S is the central atom)

7.43

■ Write Lewis structures for these molecules: (a) Tetrafluoroethylene, C2F4, the molecule from which Teflon is made, (b) Acrylonitrile, CH2CHCN, the molecule from which Orlon® is made

7.25 Define the term lone pair. 7.26

■ Based

7.28 How is electronegativity defined?

7.44 Why is it impossible for hydrogen to be the central atom in the Lewis structure of a polyatomic molecule?

7.29 Distinguish between electron affinity and electronegativity.

7.45 In the context of Lewis structures, what is resonance?

7.30 Certain elements in the periodic table shown in Figure 7.8 had no electronegativity value defined. Based on the definition of electronegativity and the identity of these elements, hypothesize as to why they have no electronegativity value.

7.46

7.31 When two atoms with different electronegativities form a covalent bond, what does the electron distribution in the bond look like?

7.48 Draw the Lewis structure of two resonance forms for ozone, O3. What prediction can you make about the relative lengths of the O!O bonds?

7.32 The bond in HF is said to be polar. For this to be true, the hydrogen atom must have less than one electron around it. Yet, the Lewis dot structure of HF attributes two electrons to hydrogen. Draw a picture of the electron density distribution for HF and use it to describe how the hydrogen atom can carry a partial positive charge. How can these two models of the HF bond (the electron density and the Lewis structure) seem so different and yet describe the same thing?

7.49 How does the existence of resonance structures for a molecule such as benzene explain its unusually unreactive chemical behavior?

7.33 Why is a bond between two atoms with different electronegativities called a polar bond?

■


Draw resonance structures for (a) NO2–, (b) NNO, and (c) HCO2–. ■

7.47 How does the structure of a molecule that exhibits resonance justify the use of the term resonance hybrid?

7.50

Consider the nitrogen–oxygen bond lengths in NO2+, NO2–, and NO3–. In which ion is the bond predicted to be longest? In which is it predicted to be the shortest? Explain briefly. ■

7.51 Which of the species listed has a Lewis structure with only one lone pair of electrons? F2, CO2– , CH4, PH3 3


237

7.52 Identify what is incorrect in the Lewis structures shown for BBr3 and SO2.

Br (a) Br

B

Br

(b) O

S

O

7.53 Identify what is incorrect in the Lewis structures shown for O3 and XeF4.

F

F

F

Xe (a) O

O

O

(b)

Shapes of Molecules 7.69 What physical concept forms the premise of VSEPR theory? 7.70

■ Predict the geometry of the following species: (a) SO2, (b) BeCl2, (c) SeCl4, (d) PCl5

7.71

■ Predict the shape of each of the following molecules or ions: (a) IF3, (b) ClO3–, (c) TeF4, (d) XeO4

7.72

■ Predict the shape of each of the following molecules or ions: (a) PH4+, (b) OSF4, (c) ClO2–, (d) I3–

7.73

Which of these molecules would be linear? Which have lone pairs around the central atom? (a) XeF2, (b) CO2, (c) BeF2, (d) OF2

7.74

■ Give approximate values for the indicated bond angles: (a) Cl!S!Cl in SCl2, (b) N!N!O in N2O, (c) bond angles 1, 2, and 3 in vinyl alcohol.

F

7.54 Chemical species are said to be isoelectronic if they have the same Lewis structure (regardless of charge). Consider these ions and write a Lewis structure for a neutral molecule that is isoelectronic with them. (a) CN–, (b) NH4+, (c) CO2– 3

■

Orbital Overlap and Chemical Bonding 7.55 Explain the concept of wave interference in your own words. 7.56 Distinguish between constructive and destructive interference. 7.57 How is the concept of orbital overlap related to the wave nature of electrons? 7.58 How does overlap explain the buildup of electron density between nuclei in a chemical bond? 7.59 How do sigma and pi bonds differ? How are they similar? 7.60

CO2, CH3OH, and CO32– all contain carbon–oxygen bonds. Draw Lewis structures for these molecules and ions. Given that double bonds are stronger than single bonds and triple bonds are stronger than double bonds, rank the four species in order of increasing C!O bond strength. ■ CO,

7.61 Draw the Lewis dot structure of the following species and identify the number of pi bonds in each: (a) CS 2, (b) CH3Cl, (c) NO2–, (d) SO2 7.62

■ Draw the Lewis dot structures of the following compounds and identify the number of pi bonds in each: (a) Cl2O, (b) H2CCH2, (c) HCCCN, (d) SiO2

Hybrid Orbitals 7.63 What observation about molecules compels us to consider the existence of hybridization of atomic orbitals? 7.64 The number of hybrid orbitals formed must always be equal to the number of atomic orbitals combined. Suggest a reason why this must be true.

1

H2 H

3

H

C

O

C

H

7.75 Propene has the chemical formula H 2C"CH!CH 3. Describe the overall shape of the molecule by considering the geometry around each carbon atom. 7.76 Describe why a central atom with four bonding pairs and one lone pair will have the lone pair in an equatorial rather than an axial location. 7.77 Describe what happens to the shape about the carbon atoms when a C"C double bond undergoes an addition reaction in which it is converted into a C!C single bond. INSIGHT INTO Molecular Scale Engineering For Drug Delivery 7.78 Describe why treatments like chemotherapy for cancer patient often have serious side effects. What is a strategy that could lessen such side effects? 7.79 Use the fundamental structure of mesoporous silica nanoparticles (MSN) to explain why they have such a large surface area. 7.80 How does a MSN differ from amorphous silica so that it has improved biocompatibility? 7.81 Draw a schematic picture of a MSN being used as a molecular engineering design selective drug delivery vehicle. Label all key components of the design. Additional Problems

7.65 What type of hybrid orbital is generated by combining the valence s orbital and all three valence p orbitals of an atom? How many hybrid orbitals result?

7.82 In 1999, chemists working for the Air Force Research Laboratory reported that they had produced a compound containing an unusual ion with the formula N5+. This ion had not been expected to exist, and its instability makes it both difficult and dangerous to prepare.

7.66 Considering only s and p atomic orbitals, list all the possible types of hybrid orbitals that can be used in the formation of (a) single bonds, (b) double bonds, and (c) triple bonds.

Consider the possible Lewis structure below. Indicate the hybridization expected for each nitrogen atom and the expected bond angles.

7.67

7.68

238

What hybrid orbitals would be expected for the central atom in each of the following molecules or ions? (a) GeCl4, (b) PBr3, (c) BeF2, (d) SO2

N

■

What type of hybridization would you expect for the carbon atom in each of the following species? (a) CO, (b) CO2, (c) H2CO, (d) CH2F2

N

N

N

Assuming that the structure shown above is correct, how many of the five nitrogen atoms would always lie in the same plane?

■


N

7.83

■ The following molecules have similar formulas, but each has a different shape. Draw Lewis structures for these

■


different electronegativities, CO2 cannot easily be separated from air by applying a voltage in the smokestack of power plant where it is generated. How does the geometry of carbon dioxide account for this observation?

compounds and then determine the expected hybridization of the central atom and the molecular shape. (a) SiF4, (b) KrF4, (c) SeF4 7.84 Consider the Lewis structure below. Is this an ion? If so, what is its charge?

O

I

O

O 7.85

A Lewis structure for the oxalate ion is shown below. (One or more other resonance forms are also possible.)

■

O

O

O

C

C

O

What is the correct charge on the oxalate ion? What type of orbital hybridization is expected for each of the carbon atoms in this structure? How many sigma bonds and how many pi bonds does the structure contain? 7.86 Methyl cyanoacrylate, C5H5NO2, is the compound commonly sold as super glue. The glue works through a process called polymerization, in which molecules of methyl cyanoacrylate form strong chemical bonds to one another and to the surfaces being glued. This process is initiated by traces of water on the surfaces. The skeletal arrangement of the atoms in methyl cyanoacrylate is shown below. Complete the Lewis structure of the compound by adding lone pairs of electrons or multiple bonds where appropriate.

H H

C

O O

C

C

H

C

N

C H

H

Using your Lewis structure, predict the hybridization for the six central atoms (the five carbon atoms and the leftmost oxygen atom). How many sigma bonds and how many pi bonds does your structure contain? 7.87 An unknown metal M forms a chloride with the formula MCl3. This chloride compound was examined and found to have a trigonal pyramidal shape. Draw a Lewis structure for MCl3 that is consistent with this molecular geometry. Use your structure to propose an identity for the metal M and explain how you have made your choice.

FOCUS ON PROBLEM SOLVING EXERCISES 7.88 Electrical engineers often use lithium niobate, LiNbO3, in designs for surface acoustic wave filters in devices such as cellular phones. What type(s) of bonding would you expect to be present in this compound? 7.89 One of the “grand challenges” for engineering recently identified by the U.S. National Academy of Engineering is the capture of carbon dioxide produced by the burning of fossil fuels. Although carbon and oxygen have significantly

■


7.90 Another “grand challenge” for engineering is managing the nitrogen cycle. Use the nature of the chemical bond in diatomic nitrogen to explain why any management is needed for this biogeochemical cycle. With so much nitrogen available in the atmosphere, why are engineering solutions needed to manage nitrogen for agricultural uses? 7.91 Lead selenide nanocrystals may provide a breakthrough in the engineering of solar panels to be efficient enough to be an economical source of electricity. Selenium is generally considered a nonmetal while lead is considered a metal. Is this distinction enough to suggest that this compound should be ionic? Explain your answer. 7.92 Both ozone, O3, and oxygen, O2, absorb UV light in the upper atmosphere. The ozone, however, absorbs slightly longer wavelengths of UV light!wavelengths that would otherwise reach the surface of the planet. (a) How can Lewis structures help explain the bonding difference between ozone and oxygen that gives rise to the difference in the absorption wavelength? (b) What relationship or relationships between light and energy do you need to know to understand this photochemical application of Lewis structures? 7.93 Nitrogen is the primary component of our atmosphere. It is also used as an inert reagent to fill containers of chemicals that might react with the oxygen in air. Draw a Lewis structure of nitrogen and use this drawing to help explain why nitrogen does not react readily with other molecules. 7.94

■ Hydrogen azide, HN 3 , is a liquid that explodes violently when subjected to shock. In the HN 3 molecule, one nitrogen–nitrogen bond length has been measured at 112 pm and the other at 124 pm. Draw a Lewis structure that accounts for these observations. Explain how your structure reflects these experimental facts.

7.95 If leads are attached to the opposite sides of a crystal of sodium chloride, it does not conduct electricity. Distilled water does not conduct electricity either. Yet, if the sodium chloride crystal is dissolved in distilled water, the resulting solution does conduct electricity. Use pictures at the molecular level of detail to describe why this occurs. Cumulative Problems 7.96 How does electronegativity relate to the periodic trends of atomic properties? How does effective nuclear charge, the concept invoked to understand those trends, also help explain the trends in electronegativity? 7.97 How do the Lewis symbols for C, Si, and Ge reflect the similarity in their electron configurations? 7.98 Use Lewis symbols to describe how free radical polymerization occurs. 7.99 When free radical polymerization occurs, how does the hybridization change the carbon atoms involved?


Copyright 2009 Cengage Learning, Inc. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part.

239

8

Molecules and Materials

OUTLINE 8.1

INSIGHT INTO Carbon

8.2

Condensed Phases—Solids

8.3

Bonding in Solids: Metals, Insulators, and Semiconductors

8.4

Intermolecular Forces

8.5

Condensed Phases—Liquids

8.6

Polymers

8.7

INSIGHT INTO The Invention of New Materials

TK The drawing shows the crystal structure for bimetallic FePt. Nanoscale particles of this material are magnetic and hold promise for extremely high density information storage devices.

T

he chemical bond is clearly one of the unifying concepts in chemistry, and our understanding of chemical bonds continues to improve. For as long as we have known that atoms join together to form molecules, scientists have continually refined the models we use to describe bonding. Active fields in modern chemical research include both computational modeling of bonds and experimental efforts to manipulate individual bonds highly selectively. Chemical bonding is the “glue” that holds molecules together and also the concept that holds together the study of chemistry. The concepts of chemical bonding that we have introduced so far may seem rather detached from the properties of the materials used in engineering designs. To bridge the gap between isolated molecules and those materials, our vantage point must be extended. In this chapter, we will look at materials such as polymers and metals and see how ideas related to chemical bonding can be applied to understand their important engineering properties. We begin by considering the various forms of carbon, an element that plays a role in many of these materials.

Chapter Objectives

Online homework for this chapter may be assigned in OWL. 240

After mastering this chapter, you should be able to ❚ describe the structures of graphite and diamond and explain how the properties of each substance arise from its structure. ❚ describe the arrangement of atoms in the common cubic crystal lattices and calculate the packing efficiency for a lattice.

❚ use band theory to describe bonding in solids. ❚ draw band diagrams for metals, insulators, and semiconductors (including n- and p-type materials). ❚ identify a material as a metal, insulator, or semiconductor from its band diagram. ❚ explain how the electrical properties of metals, insulators, and semiconductors are related to their chemical bonding. ❚ identify the types of intermolecular forces likely to be most important for a particular substance. ❚ explain the connection between intermolecular forces and properties such as boiling point and vapor pressure. ❚ describe the growth of polymers through addition and condensation reactions and predict which of these processes is likely to be important for a given monomer. ❚ describe the connection between polymer properties and molecular structure.

INSIGHT INTO

8.1

Carbon

When you think about the elements in the periodic table, you probably assume that most things are known about them. For decades, chemistry textbooks said that there were two forms of the element carbon: graphite and diamond. In 1985, that picture changed overnight. A team of chemists at Rice University discovered a new form of carbon, whose 60 atoms formed a framework that looks like a tiny soccer ball. Because the structure resembled the geodesic domes popularized by the architect Buckminster Fuller, it was given the whimsical name buckminsterfullerene. The discovery of C60 and other related molecules, now collectively known as the fullerenes, helped usher in a new form of carbon and also a new branch of science—nanotechnology. New carbon structures are certainly not the only area where nanotechnology research is flourishing, but the story of C60 is an excellent example of the way the study of chemistry affects the development of new materials. More traditional forms of carbon have long been used as materials in important applications as well. The physical properties of diamond, for example, have led to a wide range of uses. Among those properties, the impressive appearance of cut diamonds is responsible for their most familiar use in jewelry. From an engineering perspective, however, the hardness of the diamond probably represents a more significant property. Diamond drill bits find uses in many areas, including the drilling of oil wells in some locations. (The choice of drill bits depends on the nature of the rock that must be traversed.) The hardness of diamonds is sufficiently useful that industrial processes to convert graphite into small diamonds were sought for many years. Although a number of processes for making synthetic diamonds for certain applications are now viable, their use remains rather limited. Considering the cost of naturally occurring diamonds, this immediately suggests that making diamonds must be both difficult and expensive. Figure 8.1 indicates why it is so hard to make a diamond. The figure shows what is known as a phase diagram for carbon. It is a map showing which state of the element will be most stable at a given combination of temperature and pressure. If you pay close attention to the units on the vertical scale, you will see that diamond becomes the preferred form of the element only at extraordinarily high pressures—higher than 200,000 atm! Fortunately the process by which diamond would be converted to graphite is so slow as to be essentially nonexistent, so at least the diagram does not imply that our existing diamonds will soon be turned to graphite. But it does tell us that to convert bulk graphite into diamond, we will need to achieve very high pressures. (As it turns out, high temperatures are also needed to speed up the conversion process to a useful rate.) Generating and maintaining these high temperatures and pressures is difficult, expensive, and dangerous, and that’s why we don’t see synthetic

The 1996 Nobel prize in chemistry was awarded to Professors Curl, Kroto, and Smalley for the discovery of C60.

Diamond films and wafers can also be grown at very low pressure by a process called chemical vapor deposition (CVD).

8.1

Carbon

241

Figure 8.1 ❚ This phase diagram 1000 Pressure (in thousands of atmospheres)

for carbon shows which form of the element is most stable as a function of temperature and pressure. Note that pressure is shown in units of thousands of atmospheres. It is clear that graphite is the favored form at all ordinary combinations of T and P and that the formation of diamonds requires extremely high pressure.

Liquid 800

600

Diamond

400

200 Graphite 0

1000

2000 Temperature (K)

3000

Figure 8.2 ❚ Nanotubes like the one pictured here are the most recently discovered form of the element carbon. The structure of the tubes has similarities to both graphite and buckminsterfullerene. Wide ranging applications of these tubes are already beginning to emerge.

242

Chapter 8 Molecules and Materials

Image provided by Prof. R. Bruce Weisman of Rice University

diamonds in every jewelry store. As we explore bonding in solids, we will want to look into some fundamental questions about why these extreme conditions might be needed to form diamonds. How is the form of an element related to variables such as pressure and temperature? How do graphite and diamond differ in the first place so that an observable transformation from one to the other can be affected? Of course, graphite is not just a raw material for the production of diamonds; it has important applications of its own. If you happen to be writing notes in the margin of this book using a pencil, you are using graphite. What are the properties of graphite that allow it to be useful in a lead pencil, and how does the bonding in graphite give rise to those properties? Graphite also is used as a lubricant. How do lubricants work, and what is their impact on engineering designs? Other modern materials also rely on graphite as a key, if not exclusive, component. Various carbon fiber materials offer an appealing combination of high strength and low weight and are increasingly popular in sports equipment. Graphite shafts on golf clubs or carbon fibers in snow skis confirm the durability and flexibility of this material. It’s difficult to imagine making a golf club out of a bundle of pencil leads, so what is required to convert the graphite into a material that can be used in this application? When we have a composite material such as carbon fibers, what decisions must be made to control the properties? Moving from the practical to the possible, what do the new discoveries about C60 hold for materials that engineers might use in the 21st century? Although the small sphere of C60 may not have any immediate uses, manipulation of experimental conditions allows the growth of tubes of carbon, called nanotubes, as shown in Figure 8.2. These materials are not yet viable for large-scale design projects, but they

have remarkable properties. Their tensile strength is significantly higher than that of steel, and futurists have already begun to contemplate their potential for things as speculative as space elevators. Carbon nanotubes grown with a metallic element enclosed within them might represent a way to build wires that are one molecule wide. The implications for miniaturization of electronic devices are among the most immediate and exciting opportunities promised by new nanotechnologies. So scientists and engineers are actively exploring fundamental questions related to nanotechnology. How does the inclusion of metal atoms within a nanotube provide properties that might lead to a molecular wire? This question points toward the need to understand the nature of bonding in bulk samples of metals as well. The connections between molecular level characteristics and the macroscopic properties of materials are not always easy to discern, but current research in materials science and computer modeling are advancing our ability to make them. In this chapter, we will introduce ideas that can be used to infer the molecular scale explanations of why materials behave the way they do. Along the way, we will be able to answer at least some of the questions we have raised about carbon-based materials and the emerging field of nanotechnology.

8.2

Condensed Phases—Solids

The various forms of carbon that we have just described share one important physical property—they are all solids at normal laboratory temperature and pressure. The vast majority of elements are solids under such conditions. What factors contribute to the stability of condensed phases? Forces between atoms and molecules certainly play a role in the answer to this question, but the basic structure of condensed phases also contributes to an understanding of their stability. We will begin our treatment from this structural perspective. Back in Chapter 1 in our first description of the microscopic view of matter, we noted that the atoms or molecules in a solid or liquid are packed together much more closely than they would be in a gas. When we consider how those atoms or molecules are arranged into a solid, there are two broad categories. Many substances assume regular, repeating geometric arrangements and are referred to as crystalline structures. Other substances solidify into random arrangements, known as amorphous structures. Though both categories include important examples, much more is known about crystalline materials simply because their regular periodic arrays can be studied much more easily with available experimental tools. We will, therefore, focus on crystalline solids in our description here. Because atoms can be thought of as spheres, arranging them into a crystalline solid is loosely analogous to stacking a collection of balls into a box. It is intuitively clear that no matter how we arrange the balls in the box, there will always be some gaps between them; the spheres can never completely fill the volume of the box. But it should also be clear that the amount of empty space among the balls could be minimized by carefully arranging them. We can illustrate the differences in the way the balls—or atoms—may be packed by using a two-dimensional model, as shown in Figure 8.3 using marbles. It is easy to see that the gaps can be reduced if the rows of marbles are offset from one another. This idea can be quantified by thinking about the packing efficiency of a structure, which represents the percentage of space that is occupied in a given arrangement. We’ll save the calculation of this result for a problem at the end of the chapter, but the offset structure has a packing efficiency of 91%, whereas the structure in which the rows all line up has a packing efficiency of 76%. Now let’s shift our thinking from marbles to atoms. The packing efficiency of the atoms will clearly be related to the density of a material because an increase in the packing efficiency will put more atoms into the same volume. Experiments on crystal structures show that many elements condense into solids that fill space as completely

The term condensed phases refers to both solids and liquids.

Nonspherical molecules can also be packed into crystals, but we will limit our discussion to spherical atoms for the sake of simplicity.


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Figure 8.3 ❚ The marbles in the picture on the left are arranged in a very orderly pattern, but significant gaps remain. If our aim were to place as many marbles as possible into a given area, we would do better by staggering the rows, as shown in the right-hand panel. Both pictures are the same size, but the arrangement on the right allows more marbles to fit in that area.

Smaller gaps with rows offset from one another


Large gaps with rows lined up

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as possible. Such arrangements are called close-packed structures and are the threedimensional equivalent of the staggered rows of marbles in Figure 8.3. In moving from circles in two dimensions to spheres in three dimensions, though, we’ll have some additional choices to make. Let’s continue with the analogy of stacking marbles or balls into a box and suppose that we have completed a single layer of balls in the close-packed arrangement of Figure 8.3(b). When a second layer of balls is added (Figure 8.4), we could place them directly above the balls in the first layer, or we could again offset them so that each ball rests in a gap between balls in the lower

Figure 8.4 ❚ Stacking of spheres offers two ways to achieve maximum packing density. The structure on the left is called cubic close-packing, and that on the right is called hexagonal close-packing. The two arrangements differ in the way that the layers line up with one another. The cubic structure is said to feature an “a,b,c,a,b,c . . .” stacking pattern, whereas the hexagonal structure has an “a,b,a,b, . . .” pattern.

First layer of spheres arranged with alternate rows shifted with respect to one another

Second layer offers two different sets of “hollows” where we could place the third layer

Placing third-layer spheres over hollows between first-layer spheres gives cubic close-packing

You can “build” this figure using any coins you have if you want to have additional insight into the difference between the two packing arrangements.

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Stack second layer spheres into “hollows” between first layer spheres


Placing third-layer spheres directly over first-layer spheres gives hexagonal close-packing

layer. It is probably clear that this offset choice again will let us achieve higher packing efficiency and fit more balls into the box. When the second layer of balls is complete and we contemplate where to place a third layer, we face an additional choice. Again, we’ll want to offset the third layer from the second, but now there are two possible ways of doing this. The third layer of balls could be placed directly above those in the first layer, or they could be placed above hollows between the first layer balls. If the third layer is directly over the first layer, the structure is said to have hexagonal closepacking (hcp), whereas the offset third layer results in cubic close-packing (ccp). Both of these structures occupy roughly 74% of the available space. The cubic close-packed structure is often referred to as face-centered cubic (fcc), for reasons that are apparent when we look at the structure on the right in Figure 8.5. One atom is located at each corner of the cube, and additional atoms are placed at the center of each of the six faces of the cube. (Admittedly, it is difficult to see that this is the same arrangement as we saw in Figure 8.4. The best way to see it is to build models with a set of marbles or small balls.) The atoms at the four corners of a face all “touch” the atom at the center of that face. To fit an atom into the center of the face, the edge length of the cube must be larger than the diameter of the atoms, so the atoms at the corners do not touch one another. The single cube we have shown in Figure 8.5 is known as the unit cell for the face-centered cubic structure. It is the smallest collection of atoms that displays all of the features of this structure. To move from this unit cell to the structure of an actual sample of an fcc metal, we would need to extend the crystal lattice by repeating the unit cell in all directions. Some substances do not occupy space with the level of efficiency afforded by close-packed structures. For example, some metals assume a body-centered cubic structure (designated bcc), as shown in Figure 8.5. Here again, we have atoms at the corners of a cube, but now an atom is placed at the center of the cube instead of at the center of each face. The corner atoms touch the center atom but not each other, similarly to the face-centered structure. The packing efficiency in this structure is lower; the atoms occupy only about 68% of the space. In rare cases, the central atom in this bcc structure is missing. This means that the edge length of the cube becomes equal to the diameter of one atom, allowing the corner atoms to touch one another. The resulting structure is called simple cubic (designated sc). Although it is the simplest looking cubic arrangement, it is also the least efficient way to fill space. The packing efficiency is only 52%.

Simple cubic

Body-centered cubic

As you are likely to discover in a subsequent materials science course, not all unit cells are cubic, but we will restrict our discussion to those that are.

The simple cubic lattice is also sometimes called “primitive cubic.”

Face-centered cubic

Figure 8.5 ❚ The three cubic crystal lattices are shown. In a simple cubic crystal, atoms are located at each of the corners of a cube. In a body-centered cubic crystal, an additional atom sits at the center of the cube, and in a face-centered cubic crystal, atoms are found at the center of each face of the cube. Each of these arrangements repeats throughout the crystal. All of the atoms in each of these structures are identical; different colors are used only to help you to see the different positions in the lattice.

8.2 Condensed Phases—Solids Copyright 2009 Cengage Learning, Inc. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part.

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E XAM PL E PRO BLEM

8 .1

Show that the packing efficiency of the face-centered cubic structure is actually 74%, as we have claimed above.

Strategy This is essentially a geometry problem, and it requires us to understand the location of and spacing between the atoms in the fcc lattice. To find the packing efficiency, we’ll need to compare the volume of the atoms in the fcc unit cell and the volume of the cell itself. It should help if we start with a good diagram of the structure. We’ll rely in part on Figure 8.5 but will also use a simpler diagram of a single face of the cube because three-dimensional crystal structures can be challenging to draw and interpret.

Solution Before we can calculate the volume of the atoms in a unit cell, we have one very basic question to answer: How many atoms are contained in the unit cell? To answer that, we need to realize that an individual atom is frequently shared between multiple unit cells. The atoms on the faces, for example, will be in the unit cell that we are looking at as well as the adjacent one. Each of the corner atoms will appear in a total of eight unit cells because the corners of eight cubes would meet if we were to stack the unit cells up like blocks. That gives us the following result: 1 1 # atoms per unit cell = — (# face-center atoms) + — (# of corner atoms) 2 8 A cube has eight corners and six faces, giving us six face-center atoms and eight corner atoms. 1 1 # atoms per unit cell = — (6) + — (8) = 4 2 8 So each unit cell will contain four complete atoms. If we let r be the radius of an atom, then the volume of each atom is —43 pr 3. So the total volume occupied by the atoms themselves is the number of atoms multiplied by the volume of each atom. 16 4 4 atoms × q— pr 3r = —— pr 3 3 3 Now we just need to find the volume of the unit cell in terms of the radius of the atoms. Here, a drawing of a single face of the cube will come in handy. a

d = 2r

The circles represent the atoms, and the square shows the face of the unit cell cube. The diameter of an atom (d ) is 2r, and the edge length of the cube is a. From the picture, we can see that a diagonal line through the center of a face of the cube will have a length of 2d because it includes the full diameter of the center atom and the radius of each of two corner atoms. Because the face of the cube is a square, we also know that the diagonal line makes a 45º angle with the sides. So recalling a little bit of trigonometry, we can write the following relationship. a 2d

1

—— = sin 45° = —— = 0.707

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2 √‾

So a = 1.414 d = 2.828 r 3

The volume of the cube is a , and now we can express this in terms of r: Vunit cell = a3 = (2.828 r)3 = 22.63 r 3 Combine this with the volume filled by atoms from the equation above: volume of 4 atoms Packing efficiency = ————————— × 100% volume of unit cell 16

—— pr 3 16.76 3 = ————3 × 100% = ——— × 100% = 74.05%

22.63 r

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This verifies the number cited in the text.

Analyze Your Answer Because we have confirmed the result we gave earlier, it is easy to conclude that we have done the problem correctly. If we did not know the answer in advance, how might we assess our result? First, we are calculating packing efficiency, so the result must be less than 100% to be physically meaningful. The 74% efficiency of the close-packed structures is actually an upper bound. So if we did the same calculation for a different lattice, we should expect a result less than 74%.

Discussion This simple calculation helps to point out the importance of visualizing crystal structures when working with solids. Once we constructed an accurate diagram, the calculation here required only high school geometry.

Check Your Understanding Calculate the packing efficiencies for simple cubic and body-centered cubic structures. Verify your answers by comparing them with the numbers given on page 245.

Packing efficiency is not the only important factor associated with solid structures. From the perspective of bonding, an equally important concept is the coordination number—the number of atoms immediately adjacent to any given atom. Looking at the structures we have introduced thus far, we see that a simple cubic structure has a coordination number of six, a body-centered cubic structure has a coordination number of eight, and close-packed structures (both ccp and hcp) have coordination numbers of twelve. This suggests a reason that close-packed structures are much more common: by increasing the coordination numbers, they allow each atom in the lattice to interact with more nearest neighbor atoms. The attractive forces of those interactions are ultimately responsible for holding the crystal together. The periodic table in Figure 8.6 shows the preferred crystal lattice for the solid phase of each element. Scanning this table, we see that metals typically assume one of three structures: bcc, fcc, or hcp. There is no discernible pattern as to which structure is most likely for any particular metal. Nonetheless, knowledge of the structure that a metal assumes can be very important in developing new materials. Steel, for example, is predominantly iron, but by adding relatively small amounts of other elements, its properties can be fine-tuned. Deciding what atoms might usefully combine within the solid framework of iron requires that the materials engineer know what that structure is and how it might be altered by exchanging iron atoms for those of some other element. Some elements, including carbon, form more than one solid phase. (In such cases, the structure shown in Figure 8.6 represents the most stable form.) The familiar diamond and graphite forms of carbon have dramatically different physical properties, and those properties are very directly related to differences in crystal structure. The crystal structures of diamond and graphite are both rather unusual, in large part


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Figure 8.6 ❚ The shading in this periodic table indicates the most stable crystal structure for the solid state of the elements. Most of the elements display one of the structures we have discussed, but a few less common crystal arrangements are also known.

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because they are based to one degree or another on covalent bonding. The diamond structure, illustrated in Figure 8.7, is consistent with what we already know about bonding in carbon compounds. As we saw in Section 7.7, a carbon atom with four single bonds will have a tetrahedral structure. If this geometry is extended in all directions, the result is the diamond structure in which each carbon atom is bound to four neighboring atoms. Figure 8.7 emphasizes the local tetrahedral geometry for one carbon atom. The fact that each carbon atom is held in place by four strong C!C bonds explains why diamond is such a hard material. To break a diamond, we must break many covalent bonds, which would require a prohibitively large amount of energy. Silicon also displays the diamond structure. As we’ve already noted in the opening “Insight” section, though, diamond is not the only form of solid carbon, or even the most common. Graphite obviously doesn’t share the hardness of diamond, so its structure must be very different. The layered structure Figure 8.7 ❚ The carbon atoms in diamond are joined by covalent bonds, and each atom displays the same tetrahedral geometry we saw in Chapter 7 for molecules like methane. Fracturing a diamond crystal requires breaking many covalent bonds, and this explains the inordinately high strength of diamond.

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In diamond, each carbon atom forms strong covalent bonds to four other carbon atoms.


In graphite, carbon atoms form flat sheets. The atoms within a given sheet are arranged in a hexagonal pattern.

Figure 8.8 ❚ Graphite has a layered structure. The carbon atoms within each planar sheet are held together by covalent bonds, but much weaker forces hold the different sheets together. As we’ll see in Section 8.4, many of the useful properties of graphite arise from the fact that the sheets can slide across one another.

Stacking these flat sheets gives the three-dimensional structure of graphite.

of graphite (Figure 8.8) is somewhat more complicated than that of diamond. In graphite, the distance between carbon atoms within a single plane is relatively short—on the same scale as the bond length of carbon–carbon covalent bonds. Between the planes (or layers), however, the distances are much larger—more than double that within the layer. In Section 8.4, we will see that this distinction arises from fundamental differences in the forces operating within a layer versus those between layers, and it helps explain the unusual and useful physical properties of graphite. The geometric structures of solids are a very rich topic of study, and we are only able to scratch the surface of this field here. But as an engineering student, you may take an entire course later exploring the structures and properties of materials. For our present purposes, we want to introduce the structural concepts associated with solids, so that we can extend our knowledge of chemical bonding and use it to understand solids better.

8.3

Bonding in Solids: Metals, Insulators, and Semiconductors

The physical properties of metals that lead to their widespread use in countless designs can be listed readily. You are probably familiar with the basic properties of metals from science classes dating back at least to middle school. Metals are malleable, meaning they can be formed into useful shapes or foils. They are ductile, meaning they can be pulled into wires. Metals are good conductors of electricity and heat. How does the bonding in metals help explain these properties? Can the inclusion of metal atoms inside a nanotube provide the same properties? By looking at a model of metallic bonding, we can gain significant insight into these questions.

Models of Metallic Bonding In considering what we know about metallic elements, we might choose to start from the periodic trends noted in Chapter 6. Looking at metals in general and transition metals in particular, there is little difference in electronegativity from element to element. So metals and alloys are not likely to undergo ionic bonding. They are also a long way from having filled subshells, so they would require a tremendous number of covalent bonds to achieve complete octets. The experience we’ve gained so far suggests that individual metal atoms are not likely to form such large numbers of covalent bonds. If neither ionic nor covalent bonding can be expected for a metal, which models can account for the fact that metal atoms do bind to one another?

8.3 Bonding in Solids: Metals, Insulators, and Semiconductors

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We first introduced this model in Section 2.4, page 43.

The first and simplest explanation of metallic bonding is often referred to as the sea of electrons model. The essential feature of this model is that the valence electrons of metal atoms are delocalized and move freely throughout the solid rather than being tied to any specific atom, as depicted in Figure 8.9. This model accounts for several observed properties of metals. To understand malleability, consider what would happen at the atomic level if a metal were struck with a hammer. At least some of the atoms would respond to this force by moving relative to others. If the electrons are moving freely, however, as postulated in the sea of electrons model, they can adjust to the new position of the atoms and bonding is relatively unaffected. Similarly, pulling a metal into a wire changes the positions of the atoms, but delocalized electrons can adjust and maintain metallic bonding. Finally, these delocalized electrons can move in response to any applied electrical field, so the conductivity of metals can also be understood from the model. The sea of electrons model provides qualitative understanding, but quantitative models for metallic bonding also exist. The most important such model is band theory, and we can use our example of putting metal atoms into a carbon nanotube to explain the origins of this theory. If you could put lithium atoms inside a suitably sized nanotube, the metal atoms could be made to line up into a “wire” just one atom wide. (We use lithium for the simplicity of the s orbital overlap, not because this is the most practical choice.) Because the valence electron configuration of lithium is 2s1, any bonding interactions between the metal atoms will have to be based on the 2s orbitals. Those orbitals, of course, are waves, and so we can use the concept of wave interference to think about how they will interact with one another. We’ll start with the simplest case: two lithium atoms placed close enough together for their 2s orbitals to overlap one another and interact. Now recall that waves can interfere either constructively or destructively. If the pair of 2s orbitals interacts in phase with one another, the interference will be

A hammer strikes the metal, applying a force.

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Atom cores move in response to force, but electrons redistribute to maintain bonding.

Figure 8.9 ❚ The “electron sea” model of metallic bonding is somewhat crude but does account for the malleability and ductility of metals, as well as their electrical conductivity. The blue area represents the delocalized (and negatively charged) valence electrons spread throughout a piece of metal. The circles depict the positively charged cores of the metal atoms. In the left-hand picture, the atoms are arranged in a regular array. A force is applied as indicated by the large arrows—perhaps by striking the metal with a hammer. The cores of the metal atoms move in response to this force, and the sample is deformed as shown in the right-hand picture. But the free flow of electrons maintains the bonding throughout this process. A similar argument could be made for the process of drawing a metal through a die to make a wire. 250


Energy

Energy

Energy 2 Li atoms, 2 valence e–’s

4 Li atoms, 4 valence e–’s

n Li atoms, n valence e–’s

Figure 8.10 ❚ Here we examine the bonding in a one-dimensional row of lithium atoms, as might be assembled in a nanotube. In the left-hand panel, we have just two lithium atoms, each of which contributes one valence electron. With four atoms, as in the center panel, we have four electrons. As the number of atoms becomes very large, the available orbitals stack up very close together in energy. With one electron from each atom, the lower half of the available orbitals will be filled, creating the valence band. The higher energy states in the conduction band remain empty.

constructive. The resulting wave, which will have increased amplitude (or electron density) between the nuclei, is referred to as a bonding molecular orbital. On the other hand, if the pair of 2s orbitals interacts out of phase with one another, the resulting wave will have a node between the nuclei. This is known as an antibonding molecular orbital. Rather than being associated with either atom individually, these new molecular orbitals can be thought of as belonging to the pair of lithium atoms together. Figure 8.10 shows how the energies of these molecular orbitals are related to that of the original 2s orbitals. Looking at Figure 8.10, we can think through what would happen as more and more lithium atoms are loaded into the nanotube. With two atoms, there would be one bonding and one antibonding orbital. With four lithium atoms, there would be two bonding and two antibonding orbitals, as shown in the center panel of Figure 8.10. (With an odd number of atoms, one nonbonding orbital would be formed, keeping the number of bonding and antibonding orbitals equal to each other.) As we proceed further to very large numbers of atoms, the number of molecular orbitals formed becomes so large that there is virtually no energy difference from one orbital to the next. At this point, the orbitals have merged into a band of allowed energy levels, and this is the origin of the term band theory. Although our presentation here is entirely qualitative, it is important to note that a more thorough development of these ideas leads to a quantitative model. In this development, we used lithium atoms within a nanotube because doing so provided a plausible one-dimensional system. A three-dimensional bulk metal is slightly more complicated, and the way that the orbitals will combine has some dependence on the structure of the solid. For metals with valence electrons beyond the s subshell, the orbitals that ultimately form bands will include p or d orbitals. These will form additional bands, and the energy of the resulting s bands, p bands, or d bands may overlap one another. To understand properties such as conductivity, the band structure of the material provides a very powerful model.

The term “antibonding” is used because electrons in this orbital will actually weaken the bond between the atoms.

Band Theory and Conductivity The band structure of a material plays much the same role as atomic orbital energy levels did for atoms in Chapter 6. Just as the aufbau principle dictates that electrons occupy the lowest energy orbitals, electrons in bulk materials fill the bands starting at the lowest energy. The highest energy with electron occupation and the energy 8.3 Bonding in Solids: Metals, Insulators, and Semiconductors

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Insulator Metal

Semiconductor

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Conduction band

Conduction band Band gap

Valence band

Valence band

Figure 8.11 ❚ These band energy diagrams point out the key differences among metals, insulators, and semiconductors. The green shaded portions represent energy levels filled with electrons, and the blue portions depict unoccupied energy levels. Differences in conductivity can be explained in terms of the band gap and the relative difficulty of promoting electrons from filled to unfilled levels.

The energy below which an energy level is likely to be filled is called the Fermi level. For a metal, the Fermi level is at the top of the valence band.

gap between this energy level and the next available band provide a means for understanding the electrical conductivity of a material. Electrons carry a current by moving through a material, and this motion can be thought of in terms of electrons moving from one orbital to another. But electrons within a filled band cannot move readily to conduct electricity. To have mobile electrons, we will need to have an unfilled band or portion of a band that is close in energy to a filled band or portion of a band. Let’s consider these ideas for the band diagram for lithium shown on the righthand side of Figure 8.10. The band populated by valence electrons is referred to as the valence band, and the unoccupied band above it is known as the conduction band. The conduction band lies directly above the valence band in energy, and the highest occupied energy level sits at the boundary between these two bands. So an electron sitting at the top of the valence band needs to gain only a very small amount of energy to move into the conduction band. This means that lithium should be a good electrical conductor, which it is. Figure 8.11 shows some general features of the bands for three different classes of materials: metals, semiconductors, and insulators. For metals, we see the situation we just described in which addition of even an infinitesimal amount of energy can promote an electron into the valence band. The nonmetal, on the other hand, has a large energy gap between the filled valence band and the empty conduction band. So it takes a significant amount of energy to reach that empty band, where electron movement would be allowed. Under ordinary circumstances, electrons will not have enough energy to overcome this large band gap, so nonmetals do not conduct and are referred to as insulators. The diamond form of carbon is an excellent example of a nonconductive material. The intermediate case is a semiconductor, shown in the center panel of the figure. The highest filled energy level is at the top of the valence band, but the next band is only modestly higher in energy. In semiconductors, some electrons have enough thermal energy to reach the upper band even at room temperature, and therefore they have at least limited electrical conductivity. If the temperature is increased, more electrons should occupy the upper band, and the conductivity of semiconductors increases at higher temperatures. Silicon, from the same group as carbon, is the prototypical example of a pure element that behaves as a semiconductor.

Semiconductors As an engineering student, you are undoubtedly aware of the pervasive importance of semiconductors in our 21st century lives. Semiconductor technology is at the heart of the electronics industry. Now that microprocessors have found their way into our cars and appliances, the scope of that industry has never been broader. The rise of

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view of chemistry and the macroscopic scale thinking of engineering. It is fair to say that the semiconductor revolution could not have taken place without the ingenious development of a host of new materials not found in nature. Pure silicon provides an interesting textbook example of a semiconductor, but it is not useful in most applications. To generate significant conductivity in pure silicon, somehow we must promote electrons from the valence band into the conduction band. (This is the origin of the term conduction band.) This means that we must add energy, and the two most likely ways to do that are by exposing the silicon to significant heat or by allowing it to absorb light with sufficient photon energy to excite the electrons across the band gap. Neither of those would be easy to achieve or sustain in most electronic devices. Looking elsewhere in the periodic table doesn’t really offer better options, though. There are relatively few elements that can be classified as semiconductors, and none of them is suitable for widespread use in its pure form. Instead, the semiconductor industry relies on the idea that by cleverly combining elements, we can produce materials with the specific semiconducting properties desirable for a particular application. One way to do this is through a process called doping in which carefully controlled trace amounts of another element are incorporated into a semiconductor like silicon. To understand how doping works, let’s start with the band diagram of silicon shown on the left-hand side of Figure 8.12. Suppose that we wanted to increase the conductivity of pure silicon slightly. From what we have learned about the conductivity of metals, we might conclude that to do this we might put some electrons into the conduction band. So if we had some way to “inject” an extra electron, it might go into the upper band. Then it could move around in that band. If we could add enough electrons, then we might get a good conductor. How might we go about adding electrons to silicon? Each silicon atom has four valence electrons. To increase the number of valence electrons, we might try adding a small amount of an element with more than four valence electrons. Suppose we took a sample of pure solid silicon and somehow replaced one silicon atom with a phosphorus atom. Because we changed only one atom, we could assume that the overall structure of the crystal and its band diagram shouldn’t be disrupted. But now we do have one extra electron. Because the valence band is full, that additional electron must go into a higher energy orbital. Because this orbital is associated with the phosphorus atom, rather than with silicon, it will have slightly lower energy than that of the silicon conduction band. We say that the presence of the phosphorus atom introduces a donor level that lies close to the energy of the conduction band. Because the gap between this donor level and the conduction band is very small, the electron from the donor level can be promoted into the conduction band fairly easily. So in this thought experiment, adding one phosphorus atom into our silicon crystal will boost its conductivity ever so slightly.

Energy

Pure silicon

Donor electrons are promoted easily into conduction band.

Doping a material to derive useful properties is not limited to semiconductors. Many alloys, including specialty steels, are created by intentional doping.

n-type doped silicon

Donor energy level from dopant Valence band is full in pure Si.

Figure 8.12 ❚ Imagine that we add electrons to pure silicon by incorporating a trace of phosphorus. The added electrons will occupy a donor level just below the conduction band. This leads to an n-type semiconductor. The donor electrons do not need to cross the band gap to be promoted into unfilled levels of the conduction band, so they boost conductivity. We can control the number of these donor electrons by adjusting the extent of doping.


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p-type doped silicon Acceptor energy level from dopant

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Figure 8.13 ❚ In a p-type semiconductor, the addition of electron-deficient dopant atoms provides an acceptor level just above the top of the valence band. Electrons can be promoted into this level, leaving behind vacancies or “holes” in the valence band.

Valence electrons are promoted easily into acceptor level. Valence band is full in pure Si.

For a macroscopic analogy of the importance of size, think of a pile of oranges at a grocery store. If you were to replace an orange near the bottom and still maintain the pile, which would be easier to use—a tangerine (about the same size) or a large grapefruit?

It is not possible to carry out such a single atom replacement, of course. Even if it were, the resulting change in conductivity would be immeasurably small. But the same idea will hold true if we add more phosphorus atoms. The fifth valence electron from each phosphorus atom will go into a new band similar to the donor level we introduced above. Then, the donor level electrons can be promoted into the conduction band, leading to a band diagram like that shown on the right-hand side of Figure 8.12. The donor level of the phosphorus atoms provides a source of electron density near the conduction band, increasing conductivity. Moreover, the extent of that conductivity can be adjusted by controlling the amount of phosphorus present. This makes what is known as an n-type semiconductor. The name comes from the fact that we have added additional electrons, each of which has a negative charge. (Note that the overall lattice remains neutral because each phosphorus atom also has an extra proton compared to the silicon atom it replaces.) Those negatively charged electrons allow the semiconductor to carry a current. A typical n-type semiconductor might contain on the order of 0.00001% phosphorus added to a silicon crystal. (Note that to be able to do this, the silicon itself must be incredibly pure.) We say that the silicon is doped with phosphorus, and this doping usually takes place as the material is grown. As long as the dopant atoms are randomly dispersed throughout the solid, the overall effect on the crystal structure will be minimal. Other elements with more than four valence electrons could also be used. But phosphorus is a convenient choice because it has an atomic radius close to that of silicon, so it won’t disrupt the crystal lattice appreciably. The second common category of doped semiconductors is called the p-type. The idea here is to provide a way to decrease the number of electrons in the valence band, so that conductivity can be achieved without needing to promote electrons across the band gap. The strategy is similar to what we described for the n-type, except that now we want to use a dopant with fewer than four valence electrons. This introduces an acceptor level that is slightly higher in energy than the top of the valence band. The most common choice for the dopant in this case is aluminum, and Figure 8.13 shows how the band diagram will change as the material is doped. Electrons from the valence band can easily be promoted into the acceptor level, leaving behind vacancies or “holes” in the valence band. These holes behave as if they had a positive charge, so electrical engineers refer to this as a p-type material. The extent of doping controls the level of conductivity achieved, just as for the n-type.

E XAM PL E PRO BLEM

8 .2

Although much less common than silicon devices, germanium-based semiconductors can also be fabricated. Which kind of material (n- or p-type) would result if pure germanium were doped with (a) gallium, (b) arsenic, or (c) phosphorus?

Strategy Like silicon, germanium has four valence electrons per atom. So a dopant with fewer than four valence electrons will produce a p-type material, whereas a dopant with more than four valence electrons will give an n-type material. We can use the periodic table to find the number of valence electrons in each element.

254


Solution (a) Gallium is in Group 13, so it has three valence electrons. Germanium doped with gallium will be a p-type semiconductor. (b) Arsenic is in Group 15, so it has five valence electrons. Germanium doped with arsenic will be an n-type semiconductor. (c) Phosphorus is also Group 15, so germanium doped with phosphorus will be an n-type semiconductor, just as in (b).

Check Your Understanding A materials engineer needs to have an n-type semiconductor made from silicon. Which atoms might be used as dopants? Sketch a band diagram for the resulting material.

From what we’ve said so far, it probably isn’t clear why we need both types of materials or how we might make any useful devices out of them. In practice, all important circuit functions in a chip take place at the junctions between different types of materials. The point at which a piece of p-type material meets a piece of n-type material is called a p-n junction. The flow of electrons across this junction can be regulated easily by applying voltage from a battery or other power supply. Clusters of these junctions can be made to function as switches and gates of varying degrees of complexity. We won’t attempt to explain much of this here, but we can describe qualitatively how a simple p-n junction works. Figure 8.14 shows a diagram of a p-n junction. The p-type material on the left is shown with some positive holes drawn. Remember these are just missing electrons. The n-type material on the right has its extra electrons shown. Now imagine that we connect a battery across the ends of the joined materials, as shown in part (b) of the figure. The electrons from the p-type material will drift toward the positive side of the battery, whereas the holes will move toward the negative side. (Holes act as if they have positive charges.) This means that current flows across the junction when the battery is connected in this way. But suppose that we connect the voltage in the other direction, as in part (c) of the figure. Now the electrons and holes are both pulled away from the junction, so no current flows across the boundary. Our junction can serve as a volt-

Figure 8.14 ❚ A p-n junction can

⫹

⫹ ⫹

⭈⫺

⫹

⭈⫺ ⭈⫺

⭈⫺

(a) Applied voltage drives electrons and holes across junction, and current flows.

⫹ ⫹

⫹

⭈⫺

⭈⫺

⫹

⭈⫺ ⭈⫺

⫹

⫹ ⫹

⫹

serve as a simple switch, allowing current to pass when a voltage with the correct polarity is applied. This simple functionality can be exploited to build a variety of complex circuit functions.

⫺

⫹

(b)

⫺

For a common type of device—the metal oxide semiconductor field effect transistor (MOSFET), the voltage needed is in the range of 1–3 V.

⭈⫺ ⭈⫺

With polarity of applied voltage reversed, no current flows across junction. ⫹

⭈⫺ ⭈⫺

(c)


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age-controlled switch: we can turn the current across the junction on or off by applying the voltage in the right direction. This simple idea can be applied to develop circuits of great complexity. That takes a lot of junctions of various sorts and a lot of clever design, but it can be done. Many of you will go on to take a materials properties course in engineering, where you will probably see these ideas from a different perspective. But for our purposes here, the main aim is to show how important bulk properties can be manipulated by adjusting the atomic composition of a solid. The features of metallic conductors, semiconductors, and insulators that we’ve outlined here apply to more than just elemental forms of substances. Most alloys are combinations of metallic elements (with an occasional minor component that is a nonmetal), and their bonding resembles what we have just described for metals. Semiconductors are often made from combinations of elements, too. Gallium arsenide is a particularly important example. Insulators, too, are often made from combinations of elements, and many ceramic materials are insulators. Not all materials have chemical bonds that are responsible for holding a solid together. Now we will look at other forces that play roles in condensed matter.

8.4

Intermolecular Forces

In Chapter 7, we saw that the formation of chemical bonds is always governed by minimization of the energy of the atoms involved. As we turn our attention to the structures of solids, we’ll see that the same principle still holds. The structure of any solid, whether crystalline or amorphous, is determined by the balance of attractive and repulsive forces between the atoms or molecules involved. In this section, we will consider the nature of intermolecular forces—the forces between molecules.

Forces Between Molecules Intermolecular forces are also critically important in biochemical systems, including several of the biomaterials we considered in Chapter 7.

In Chapter 7, we learned how interactions between atoms lead to bonds that hold molecules together. There’s no doubt that those interactions are an extremely important facet of chemistry. But understanding chemical bonding still leaves us a bit short of understanding many aspects of the structure of matter. The forces between molecules are also important. Though typically much weaker than the forces in chemical bonds, these intermolecular forces are largely responsible for determining the structure and properties of condensed phases. So they are especially important from an engineering perspective. Here we will look at various types of intermolecular forces.

Dispersion Forces Dispersion forces, sometimes called London forces, are common to all molecules. They are also referred to as instantaneous dipole–induced dipole forces. This rather awkward sounding term points us to the origin of these forces, so let’s consider each part of the name separately. To begin with, recall that a dipole exists whenever we have two oppositely charged points or objects separated by some distance. The quantum mechanical picture of the atom, and hence our pictures of various orbitals, arises from ideas of probability and average position. To understand the notion of an instantaneous dipole, though, we’ll need to step away from that viewpoint. Instead, we’ll need to consider electron positions or distributions at a single instant. For an electron in an s orbital, for example, the ideas presented in Chapter 6 tell us that the probability of finding that electron at some distance from the nucleus is spherically symmetrical. But now imagine that we could take a single instantaneous snapshot of the electron. You’ll probably begin to think that the electron is more likely to be found on one side of the nucleus. (A similar snapshot a moment later

256


might show the electron on the other side, so the average position will still be consistent with the spherical shape we ascribe to the orbital.) This dynamic fluctuation in the distribution of electrons holds in molecules as well. Thus, at any given instant, we may find more electrons on one end of a molecule than on the other, even if, on average, they are evenly dispersed throughout. Whenever such a fluctuation occurs, there will be a momentary asymmetry of charge; one end of the molecule will have a small, transient negative charge, and the other will have a small, transient positive charge. This immediately implies a short-lived charge separation, which gives rise to a dipole moment. This dipole differs from those described earlier in our consideration of polar bonds because it lasts only for that instant that the electrons happen to be unevenly distributed. Accordingly, it is referred to as an instantaneous dipole. The lifetime of this instantaneous dipole is fleeting, and it soon disappears only to form again as the momentary distribution of electrons becomes asymmetrical again. This dynamic distribution of electrons is not the only way to achieve an uneven distribution of electron density. Because electrons have very little mass and are charged, they can also be influenced by electric fields external to the molecule, such as those from a dipole in a neighboring molecule. When a molecule encounters an external electric field, the negatively charged electrons will tend to migrate away from the negative pole of the applied field. This response establishes a new dipole within the molecule itself. Thus the external field has forced a dipole to exist where it had not existed previously; it has induced the dipole. This process is shown schematically in Figure 8.15. Now we combine the ideas of the instantaneous and induced dipoles. Consider two molecules that are separated by a small distance. One molecule experiences a fluctuation in its electron density, creating an instantaneous dipole. The second molecule will then experience that dipole as an external electric field, creating an induced dipole in the second molecule. This combination is the instantaneous dipole–induced dipole interaction, or dispersion. Because the dipoles paired in this way are aligned so that the positive end of one is near the negative end of the other, dispersion forces are attractive. You may be thinking that such fleeting forces between pairs of molecules must be incredibly weak—and if so, you’re right. But because macroscopic samples contain on the order of Avogadro’s number of molecules, the sum of all these weak interactions can be a significant amount of energy. This energy holds many liquids and some solids together. How does one estimate the relative size of dispersion forces for different materials? The key concept needed to answer this question is polarizability. This term indicates how susceptible a molecule is to having its charge distribution perturbed by an external field. This feature of molecules can be quantified experimentally by measuring the strengths of interactions between molecules, and there are some general conclusions that we can draw. The most important of them is that large molecules tend to be more polarizable than small ones. Because the interaction between positively charged

Many polymers, including polyethylene, form solids in which dispersion forces account for the most important intermolecular attractions.

The unperturbed molecule has a symmetric charge distribution.

δ⫹

Negatively charged electrons are repelled by the negative side of the external field. This establishes an induced dipole in the molecule.

δ⫺ δ⫹

δ⫺

Figure 8.15 ❚ An external electric field distorts the shape of the electron cloud around a molecule or atom, creating a transient or induced dipole. The distorted molecule can then act as an external field on a nearby molecule, propagating the effect. Dispersion forces result from the mutual attractions among such induced dipoles in a collection of molecules.


257

nuclei and negatively charged electrons decreases as the distance increases, it makes sense that large molecules are more polarizable. The electrons are less tightly held and thus more easily influenced by the external field. The fact that larger molecules are more polarizable means that they tend to experience stronger dispersion forces.

Dipole–Dipole Forces When molecules that have permanent dipoles interact with each other, there is a tendency for those dipoles to align. This alignment, particularly in a liquid, is not rigid. But the positive region of one molecule prefers to be near the negative end of other molecules and vice versa. The charge at these poles, however, is seldom more than a fraction of the charge of an electron. Thus, although the extent of the attraction of opposite charges or repulsion of the same charges is more than that in dispersion forces, it is quite modest compared to the strength of chemical bonds. With only small attractions and repulsions operating, the alignment of the dipoles will not be absolute. At any given time, there will be pairs or groups of molecules in a material that experience repulsive interactions, but there will be more attractive ones. Figure 8.16 shows this situation for a snapshot of a 50-molecule sample. Solid black lines indicate attractive interactions and dashed red lines show repulsive interactions. There are clearly more attractive interactions than repulsive, and this picture shows only a small number of molecules. If this snapshot were taken moments later, the location of the various interactions would change, but the number of attractive and repulsive ones would remain similar, on average. Because they are based on permanent rather than transient dipoles, these dipole–dipole interactions are typically stronger than dispersion forces.

Hydrogen Bonding The remaining intermolecular interaction is among the most poorly labeled terms in all of science. Although they are frequently called “hydrogen bonds,” hydrogen bonding interactions are not really bonds at all. They do not involve a chemical bond in the same formal sense in which we used that term in Chapter 7. They are, rather, an especially strong form of dipole–dipole interactions that are observed when certain combinations Figure 8.16 ❚ This drawing shows an arrangement of 50 polar molecules, as might be found in an instantaneous snapshot of a liquid. The solid black lines indicate attractive interactions between opposite ends of the dipoles of pairs of molecules, and the dashed red lines indicate repulsive interactions. The number and strength of the attractions outweigh those of the repulsions, so there is a net force holding the molecules together.

Attractive interactions

Repulsive interactions

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of elements are present in compounds. Hydrogen bonding is typically the strongest of all intermolecular interactions. Because hydrogen bonds play a prominent role in biological chemistry, these interactions have a profound effect on life as we know it. There are some specific criteria that must be met for hydrogen bonding interactions to form. The “hydrogen” in hydrogen bonding is always a hydrogen atom bound to a highly electronegative atom, typically N, O, or F. The highly electronegative atom will tend to attract the electron pair it shares with hydrogen, and so the hydrogen atom will take on a partial positive charge. Because hydrogen has no other valence electrons, this positive charge will be especially “exposed” to other molecules. If an atom with a partial negative charge and an available lone pair of electrons approaches this exposed positive charge, a strong dipole-based attraction will result. The list of atoms best able to provide this source of electron density is the same as mentioned before: N, O, and F. In virtually all of their compounds, these elements tend to have one or more lone pairs of electrons available, and these electrons can be at least partially shared with a nearby hydrogen atom from another molecule. The simplest cases of hydrogen bonding involve pure compounds whose molecules contain both the hydrogen atom and the highly electronegative atom with which it will interact. Thus in hydrogen fluoride, the H atom of one molecule can be attracted toward the F atom of a second molecule. A similar situation exists for water, where the oxygen atoms provide the lone pairs to which the hydrogen atoms will be attracted. It is crucial to remember that these hydrogen bonding interactions are much weaker than the chemical bonds within a single molecule. The hydrogen bonds in water, for example, are responsible for holding the molecules to one another in the solid or liquid phases. When liquid water is heated, the hydrogen bonds break down, but the chemical bonds do not. Thus as the water boils, the vapor driven off is still made up of H2O molecules. (Recall Example Problem 1.1, page 8.) In more complicated cases, including many that arise in biological molecules, hydrogen bonding can occur between different molecules or even between different parts of large molecules like proteins. The common thread in all of these instances is the presence of hydrogen atoms bound to highly electronegative atoms.


Among the many roles for hydrogen bonding, perhaps the most important is the linking of two strands of DNA into a double helix.

8.3

Which type(s) of intermolecular forces need to be overcome to convert each of the following from liquids to gases? (a) CH4, (b) CH3F, (c) CH3OH

Strategy Dispersion forces will exist for any substance but will be the most important forces only in the absence of stronger dipole–dipole or hydrogen bonding interactions. To see if those stronger forces are possible, we need to look at the structure of each compound. So we might start by using Lewis structures to establish the molecular geometries. In this case, we are dealing with methane, whose structure we know well, and two closely related compounds.

Solution (a) We are quite familiar with the structure of CH4 from Chapter 7: it is a tetrahedral molecule with four identical C!H bonds. H H

C

H

H Because the structure is totally symmetrical, no dipole–dipole forces are possible. Because the H atoms in CH4 are not bound to N, O, or F, no hydrogen bonding is possible. Only dispersion forces need to be overcome to vaporize liquid methane.


259

(b) Methyl fluoride also is tetrahedral, but one of the hydrogen atoms in methane has been replaced by fluorine. F H

C

H

H With only one fluorine atom present, the molecule is no longer symmetrical. The C—F bond is highly polar and will give the molecule a dipole moment. So to vaporize CH3F, dipole–dipole forces must be overcome. (c) The structure of methanol is also closely related to that of methane. Here we replace one of the hydrogen atoms in CH4 with an !OH group. H H

C

O

H

H As in CH3F, the molecule will have a dipole moment, but the dipole arises from the —OH group. That same functional group means that hydrogen bonding will be possible here. The hydrogen atom in the !OH group of one molecule can be strongly attracted to a lone pair on the oxygen atom in another molecule. So to vaporize methanol, we need to overcome hydrogen bonding interactions.

Discussion Identifying the forces involved lets us predict the order of boiling points for these liquids. With only weak dispersion forces involved, CH4 should have a very low boiling point. Dipole–dipole interactions should make it more difficult to vaporize CH3F, giving it a higher boiling point. (Dispersion forces would also be slightly stronger than for CH4 because of the increased mass.) Finally, hydrogen bonding in CH3OH should give it the highest boiling point among this group. These predictions are correct, as the boiling points are –164ºC (CH4), –78ºC (CH3F), and 65ºC (CH3OH).

Check Your Understanding Rank the following in order of increasing boiling points by considering the important intermolecular forces in each: Ne, CO, CH4, NH3.

The forces we’ve considered here—dispersion, dipole–dipole interactions, and hydrogen bonding—give rise to many of the interactions that can occur between molecules and allow us to understand both the existence and properties of condensed phases. None of these forces is even as strong as fairly weak chemical bonds. Hydrogen bonding interactions are typically the strongest of the intermolecular interactions, but even they are an order of magnitude weaker than the average covalent bond. Despite this apparent weakness, when considering a macroscopic sample, the sheer number of such interactions can make them a key factor in determining bulk properties. The comparison between intermolecular and intramolecular forces can be illustrated very clearly by looking more closely at graphite. (Recall the general features of the structure from Figure 8.8.) Within a plane, the carbon atoms are bound by strong covalent bonds, as shown in the left-hand panel of Figure 8.17. A large number of resonance structures could be drawn, but it is reasonably clear that, on average, the strength of each carbon–carbon bond will be between those of a single bond and a double bond. Between the sheets, however, the interactions are due to dispersion forces, making them dramatically weaker. So the carbon atoms within a given sheet are held together by relatively strong chemical bonds, whereas much weaker dispersion forces hold pairs of adjacent sheets to one another. The result is that sheets of graphite can slide past each other relatively easily and this helps explain why graphite 260


C—C bond length = 141.5 pm

Spacing between sheets = 335 pm

“Top view” of a graphite sheet

“Side view” of a stack of three graphite sheets

Figure 8.17 ❚ Carbon atoms within each sheet in the graphite structure are joined by strong covalent bonds, with a fairly short C—C distance. Because each carbon atom forms four bonds within its own sheet, intermolecular forces must be responsible for holding the sheets together. This is apparent from the much larger spacing between layers. Because the forces holding the layers to one another are weak, sliding of one layer relative to the next is fairly easy, and this explains the softness and lubricating properties of graphite.

is useful as a lubricant. Lubricants must withstand large forces, and the strong bonds within a sheet provide that strength. They must also be capable of viscous flow, and the movement of one sheet relative to the next makes that possible. Though fairly simple, the structure of graphite points to the importance of both chemical bonding and intermolecular forces in determining the properties of a solid. Next we turn our attention to the influence of intermolecular forces on the physical properties of liquids.

8.5


Solids are not the only condensed phase of matter. Intermolecular forces of the type we just described often result in the formation of liquids as well as solids. From a structural perspective, liquids are not as orderly as solids, but we can look into some of their important characteristics. The key difference between liquids and solids lies in the much greater mobility of the atoms or molecules in a liquid. For a solid, the average position of each particle remains essentially unchanged; even though there may be significant vibration about that average position, the atoms remain in a fixed arrangement over time. By contrast, particles in a liquid are free to move with respect to one another, and they do so constantly. This constant motion plays an important role in the fact that liquids often serve as solvents that can dissolve other chemicals. The dissolved particles, like those of the liquid itself, also become capable of movement. This allows dissolved particles to encounter other dissolved materials (or solids in contact with the liquid), and such encounters may allow chemical reactions to take place. Confining our discussion for the moment to pure liquids, though, what can we learn about their properties from looking at intermolecular forces? Let’s consider some physical properties that might be important in understanding the behavior of liquids.

Vapor Pressure After a springtime rain shower, the sidewalks are covered with puddles. Those puddles disappear but clearly the water never boils. The evaporation of these puddles is indicative of a phenomenon known as vapor pressure that is actually characteristic of both liquids and solids. Vapor pressure is the gas phase pressure of a substance in 8.5 Condensed Phases—Liquids

261

Vapor pressure measurmenets are made in evacuated vessels, so the only source of pressure is from the vapor.

Figure 8.18 ❚ When a liquid is held in a closed container, it will establish equilibrium with its vapor phase. The vapor pressure provides a measure of the ease with which a particular liquid evaporates.

262

equilibrium with the pure liquid in a sealed container. It is a characteristic property of a particular substance at a particular temperature. A puddle of the same size will evaporate faster on a hot day than on a cool one. If the liquid on the sidewalks were acetone rather than water, the puddles would disappear much more quickly. We can reach a qualitative description for the idea of vapor pressure by considering intermolecular interactions of the various types we introduced in the previous section. In any solid or a liquid, these attractive intermolecular interactions hold molecules together. From a molecular perspective, the system consists of a large number of molecules, constantly moving but interacting with one another. In any molecular system, the velocities of individual particles vary, but there are always some particles with high kinetic energy. (The Boltzmann distribution discussed in Section 5.6 for gas molecules is an example of this. Similar ideas apply to molecules in liquids, too, although their exact form is more complicated.) If an especially energetic molecule happens to find its way to the surface of the liquid, it may have enough energy to overcome the attractive tugs of its neighboring molecules and escape into the vapor phase. That is the molecular origin of evaporation. This molecular viewpoint lets us predict the behavior of vapor pressure with temperature changes. As temperature increases, the number of molecules with high kinetic energy also increases. So, more molecules will be capable of escaping the attractive forces of intermolecular interaction. Thus vapor pressure should increase as temperature increases, and this is seen experimentally. Our molecular description also allows us to understand the relative vapor pressures of different molecules. To escape from a liquid, molecules must have kinetic energy sufficiently high to overcome the intermolecular interactions in the liquid. So the strength of those intermolecular forces sets the bar that evaporating molecules must overcome to escape. If a system has strong intermolecular interactions, molecules will need particularly large amounts of kinetic energy. There are simply fewer molecules capable of escaping when intermolecular interactions are strong, and so the vapor pressure will be low. Liquids with high vapor pressures evaporate rapidly and are said to be volatile. Many engineering designs, including those of gasoline powered engines, must take into account the volatility of some component, such as the fuel. Volatility is sometimes an asset, for example, when fuel is being injected into a cylinder for combustion. It can also be a liability, for example when trying to minimize the release of unburned fuel from a storage tank into the atmosphere. This gives us a reasonably useful qualitative view of vapor pressure. But how can it be quantified? Water has a lower vapor pressure than acetone, but how much lower? To measure the vapor pressure of a liquid or a solid requires that the system reach equilibrium. In other words, the molecules that are escaping from the liquid must be replaced by other molecules that are being captured, as shown in Figure 8.18. Once this equilibrium has been established, the pressure of the molecules in the gas phase

A nonvolatile liquid evaporates slowly and has a low vapor pressure.


Volatile liquids evaporate more easily, resulting in higher vapor pressure.

Table

❚

8.1 Recall that 760 torr = 1 atm.

Vapor pressures at 295 K and normal boiling points of a variety of substances

Substance

Vapor Pressure (torr)

Normal Boiling Point (ºC)

Acetone

202

56.2

Br2

185

58.8

CClF3

24,940

–81.1

CCl2F2

4448

–29.8

CCl3F

717

23.8

CCl4

99.0

76.54

HCN

657

26

Formaldehyde

3525

–21

Methanol

108

64.96

n-Pentane

455

36.07

Neopentane

1163

9.5

Isobutane

2393

n-Butane

1658

–0.5

Propane

6586

–42.07

Ethane

29,380

–88.63

Water

19.8

–11

100

will remain constant and can be measured fairly easily. This process, in which the net state of the system is unchanged but there are still molecules escaping and being captured by the liquid, is called dynamic equilibrium. We’ll study this idea in detail in Chapter 12 because it is important in many more circumstances than just vapor pressure. The values of the vapor pressure for a number of liquids near room temperature are provided in Table 8.1.

Boiling Point The boiling point of a liquid also depends on the strengths of the intermolecular interactions and is closely related to vapor pressure. As just noted, the vapor pressure of a material increases as temperature increases. This phenomenon is true over all temperature ranges, so that eventually as temperature continues to increase, the vapor pressure of the liquid equals the external pressure. When that happens, small bubbles of vapor form within the liquid and the system boils. The normal boiling point of a liquid is defined as the temperature at which its vapor pressure is equal to one atmosphere. The well-known normal boiling point of water is 100ºC. If the external pressure is below one atmosphere, the temperature at which boiling will occur is lower. This idea gives rise to the special cooking instructions for people who live in the mountains, where atmospheric pressure is reduced. If one wishes to boil an egg at high altitude, it will take longer than at sea level because the water is boiling at a lower temperature. The vapor pressure of water as a function of temperature

Formation of gas bubbles in a liquid is not automatic, so it is possible to “superheat” a liquid above its boiling point under certain circumstances.

8.5 Condensed Phases—Liquids

263

Figure 8.19 ❚ The vapor pressure of water varies dramatically as a function of temperature.

Vapor pressure is 1 atm (760 torr) at the normal boiling point, which is 100°C for H2O. 1140

Vapor pressure (torr)

950

Note that vapor pressure continues to increase past the boiling point. Familiar things like popcorn or pressure cookers rely on this fact.

760 570 380 190

0

20

40

60 Temperature (°C)

80

100

120

is shown in Figure 8.19. The difference in the boiling temperature between sea level and the mountains is a direct result of the temperature dependence of vapor pressure. At higher altitude, atmospheric pressure is lower, and boiling occurs whenever the vapor pressure is equal to the external pressure. Note also the shape of the curve. The pressure increases rapidly as temperatures exceed 100ºC, where water typically boils. The normal boiling points of several liquids are listed in Table 8.1. Note that those liquids with low vapor pressures have high boiling points. It is possible therefore to rate the strength of intermolecular interactions by observing the normal boiling point of liquids: the higher the boiling point, the stronger the intermolecular forces. It is also possible to separate molecules of different intermolecular interaction strengths by distillation, which is an important step in oil refining.

Surface Tension

In engineering applications where surface tension can create problems, molecules called surfactants are often added to decrease the surface tension. Surfactants also rely on an understanding of intermolecular forces in their design.

264

The attractive nature of intermolecular interactions means that a liquid will, in general, behave to maximize the number of interactions. One effect of this is that the number of molecules at the liquid’s surface is minimized. Such surface molecules will find themselves in a different environment than those in the bulk of the liquid. If the surface of the liquid is in contact with a gas, there are fewer molecules available to interact with these surface molecules because gases are less dense than solids (Figure 8.20). Because the typical intermolecular interactions are attractive, the fact that these surface molecules have fewer neighbors means they find themselves in a state of higher energy than bulk molecules. The system will naturally minimize the number of molecules in this higher energy environment, giving rise to a phenomenon called surface tension. Something under tension is in a higher energy state. Here, the origin of the high energy state is the existence of the surface and the fact that molecules at the surface have fewer nearest neighbors than molecules in the bulk of the liquid. Surface tension explains a variety of observable phenomena, including the fact that water tends to bead up into spherical (or nearly spherical) drops. The energy of any given sample of liquid will be reduced if fewer molecules sit at the surface. The shape that encloses the greatest volume with the least amount of surface is a sphere. The fact that water experiences strong intermolecular interactions due to hydrogen bonding means that its surface tension is relatively high and provides a strong force to form spherical drops. This effect is best seen on a waxed surface because the polar water molecules do not interact strongly with the nonpolar wax. On other surfaces, including a dirty car hood, strong interactions between the water molecules and substances on the surface can overcome the surface tension, producing a more uniform film of water.


Molecules at the surface of a liquid have fewer neighbors, so they experience fewer attractive forces.

Figure 8.20 ❚ Molecules at the surface of a liquid have fewer nearby neighbors than molecules in the bulk of the liquid. Surface tension arises because of this difference in the molecular scale environment.

When the surface of a liquid interacts with a solid, rather than with a gas, we must consider competing interactions. The liquid molecules interact with each other, but they also interact with the molecules in the solid. The liquid–liquid interactions are referred to as cohesion, and the liquid–solid interactions are called adhesion. The relative strengths of these forces will dictate the shape of the curved surface, or meniscus, that forms in the liquid. Figure 8.21 shows a comparison of two different liquids in a glass tube. Water (on the top) has strong interactions with the glass, which has numerous polar bonds on its surface. So the adhesion forces are larger than cohesion and the meniscus is concave. The liquid in the bottom of the tube is mercury, which does not interact strongly with the polar glass surface. So adhesion is low relative to cohesion and a convex meniscus results. A liquid in which the adhesive and cohesive forces are equally strong would be expected to show a horizontal surface, with virtually no meniscus. The carbon structures we noted in the opening “Insight” section in this chapter do not readily exist in liquid phases. To melt diamond or graphite requires temperatures that are high enough that there are no practical consequences of the liquid phases. There are, however, carbon-based materials that have important properties in both the solid and liquid states, and they comprise a sizable portion of the materials used in modern engineering designs. Before we can move on past our study of materials, we need to take another look at polymers—particularly carbon-based polymers.

8.6

Polymers

The importance of carbon in chemistry and materials science is due to the myriad chemical compounds it can form, not to the elemental forms we have noted thus far in this chapter. Though there are more than enough organic compounds to provide a complete course devoted to their properties and reactions, we will limit our consideration here to carbon-based polymers. As we saw back in Chapter 2, polymers are giants among molecules, constructed by the sequential stringing together of smaller molecules called monomers. They constitute the largest component of the materials we call plastics—though plastics are defined by their ability to be molded and usually include additives.


Molecules in the bulk of a liquid have more neighbors, so they experience more attractive forces than do molecules at the surface.

Figure 8.21 ❚ The graduated cylinder shown in this photograph contains a layer of water on top of a layer of mercury. The meniscus at the top of the mercury column is concave down, meaning that the mercury is higher in the center of the cylinder than at the walls. The meniscus at the surface of the water is concave up, with the water higher along the walls than in the center of the cylinder. These shapes result from the differences in strength of the adhesive and cohesive interactions between the two liquids and the glass of the graduated cylinder.

8.6

Polymers

265

The utility of polymers and plastics in engineering applications arises from the fact that materials chemists and engineers can control their physical properties. The main factors that can be empirically adjusted to modify polymers are the monomers used, the type of reactions needed to generate the polymers, and the catalysts that are employed to speed the reactions. A careful choice of these factors can ultimately control the physical properties of the resulting polymer.

Addition Polymers Now that we understand a bit about chemical bonds and chemical reactions, it should not seem surprising that the manner in which polymers are generated depends strongly on the characteristics of the monomers involved. In the closing “Insight” section of Chapter 2, we showed how monomers containing one or more double bonds undergo free radical addition reactions to form polymers. As shown in Figure 8.22, an addition polymerization begins with the generation of a free radical. (This step is normally accomplished by heating a molecule that decomposes into two radicals.) This free radical then attacks the double bond in a monomer molecule, forming a new radical that now includes one monomer unit. The resulting radical attacks the double bond of another monomer, breaking it and forming an additional single bond. Importantly, the product is still a radical, now containing two monomer units. This radical

Figure 8.22 ❚ Addition polymerization often proceeds by a free radical mechanism, featuring initiation, propagation, and termination steps. Conditions are chosen so that the initiation and termination steps occur infrequently. Each propagation step adds one monomer unit to the growing polymer chain.

Initiation step produces free radicals by breaking down a molecule like this organic peroxide.

R

O

O

R

R

H R

R

H C

O

O ⫹ O

R

C

H

O

H

H

H

C

C

H

H

H

H

H

H

C

C

C

C

H

H

H

H

Propagation steps add monomers to the growing chain. Because the end of the chain remains a free radical, these steps can continue as long as monomer is present.

Peroxides contain a pair of oxygen atoms connected by a single bond.

O

R

O

O

H

H

C

C

H

H

H

H C

H

R

C

O

H

Termination steps occur when two radicals encounter one another and combine. This eliminates both radicals and stops the growth of the polymer chain. If the number of radicals present is kept low, termination steps will be rare.

R

266

O


H

H

H

H

C

C

C

C

H

H

H

H

R

O

H

H

H

H

C

C

C

C

H

H

H

H

can attack yet another monomer, and with enough of these steps, the polymer can be formed. Just how many times this reaction repeats itself controls the length of the polymer chain. Eventually, a radical will be eliminated, either by encountering another radical or by some other reaction. This stops the polymer molecule from growing longer and is referred to as chain termination. Because the occurrence of these termination steps is somewhat random, not all of the polymer chains will grow to the same length. Instead, there is normally a range of lengths and hence a distribution of molecular masses. This distribution can be roughly described in terms of the degree of polymerization, which is the average number of repeating units in the polymer. This is usually calculated in terms of mass: molar mass of polymer Degree of polymerization = ——————————— molar mass of monomer One reason that polymers are such common engineering materials is that their properties can be tailored in a number of ways. The most obvious is the choice of the monomer unit because we know that different polymers can offer vastly different characteristics. But it is also possible to have distinctly different properties in polymers derived from the same monomer. One variable is the degree of polymerization, which can be controlled by varying parameters used in the synthesis reaction. Another approach is to control the way in which the monomers are linked together. Perhaps the most important example of this is the formation of both low density polyethylene (LDPE) and high density polyethylene (HDPE), which we explored in the closing section of Chapter 2. When the monomer is less symmetrical than ethylene, there is additional variability that might result from the polymerization process. For example, propylene has a structure similar to that of ethylene but with a methyl (–CH3) group replacing one hydrogen atom. This change means that the resulting polymer may have methyl groups attached to the carbon backbone in three different ways, resulting in the various geometric arrangements shown in Figure 8.23. It is important to realize that the geometry at each carbon atom in the chain is tetrahedral, so the methyl groups are

(a)

(b)

(c)

Figure 8.23 ❚ For monomers with more complicated structures, different ways of linking them together become possible. Here we see polypropylene, which differs from polyethylene in that every other carbon atom has a methyl group in place of one hydrogen atom. The large purple balls here represent those methyl groups. In (a), all of the methyl groups are arranged on the same side of the polymer chain, giving isotactic polypropylene. In (b), the methyl groups are on alternate sides of the chain, making syndiotactic polypropylene. Finally, the random arrangement of the methyl groups in (c) is atactic polypropylene. 8.6

Polymers

267

Polystyrene is an example of a commercial polymer whose properties depend on whether it is syndiotactic or atactic.

either in front of or behind the plane of the page in this illustration. If the methyl groups all point forward (as in Figure 8.23(a)) or backward, then the polymer is said to be isotactic. If the position of the methyl groups systematically alternates between forward and backward, the configuration is called syndiotactic. In the third case, the arrangement of the methyl groups is random, sticking “into” and “out of ” the page with no regular pattern, and the polymer is said to be atactic. The physical properties of a polymer can be strongly influenced by this type of variability. Because the structure of atactic polymers is not very well defined, their physical properties can vary considerably from one sample to another. The regular arrangement of the monomers in isotactic and syndiotactic polymers generally leads to more predictable and controllable properties. So our ability to produce polymers selectively in one of these arrangements offers a powerful tool to the materials scientist or engineer. This ability to control the specific configuration of a polymer was first achieved by Karl Ziegler and Giulio Natta in the 1950s. These scientists discovered new catalysts for the addition polymerization reaction that increased the rate of reaction and also controlled the structure. The catalysts came to be known as ZieglerNatta catalysts. Their discovery invigorated the study of polymers, and they were awarded the Nobel prize in chemistry in 1963 for their efforts.

Condensation Polymers The second commonly observed type of reaction for the formation of polymers is called condensation. In this reaction, there are two functional groups on the monomer, which can react together to form small molecules. Those small molecules then split off, and the remaining parts of the two monomers are joined together. (The term condensation polymer originates from the fact that the small molecule eliminated is often water.) Figure 8.24 shows the condensation reactions that produce two polymers

Figure 8.24 ❚ In condensation polymerization, a molecule of water is eliminated as each monomer is added to the chain. Here we show the first steps in the formation of Nylon (upper panel) and Dacron (lower panel). Both are copolymers, with two different monomer molecules combining to form the polymer. Because each monomer contributes its own functional group to the condensation step, regular alternation of monomers is ensured. Other condensation polymerizations can involve a single monomer with different functional groups on each end.

H

O OH

HO

⫹

H

N

O

Adipic acid

N Hexamethylenediamine

H

O

N

HO

N

O

Amide linkage

H

O OH

HO

O

⫹

HO

Ethylene glycol

OH Terephthalic acid

O

O ⫹ H2O

HO

O

OH Ester linkage

268


H

H

⫹ H 2O

H

that are important in fiber manufacturing, Nylon® and Dacron®. Nylon is referred to as a polyamide, and Dacron is a polyester. These names refer to the characteristic amide and ester linkages found in the molecules, as highlighted in Figure 8.24. Looking at these two examples, we can see the functional groups that must be present in monomers for condensation polymerization.


8.4

Each of the following monomers or pairs of monomers can undergo condensation polymerization reactions. Draw the structures showing the repeat units and linkages in each of the resulting polymers. (a) glycine: OH

H2N O (b) 6-aminocaproic acid:

O H2N

OH

(c) p-phenylenediamine and terephthalic acid: O H2N

O

NH2 HO

OH

Strategy Condensation polymerization involves the elimination of a small molecule—usually water—as each monomer is added to the growing chain. We can look for functional groups on the ends of the monomers that can undergo such a reaction.

Solution (a) We have an amine group (!NH2) on one end of the monomer and a carboxyl group (!COOH) on the other. Elimination of H2O from these groups will produce an amide linkage. The polymer will have the following structure, with the monomers connected through amide linkages: H

O

[N

]

N H

O

(b) This is similar to (a). H

[N

O

]

N H

O

8.6

Polymers

269

(c) Here, we have amide groups on one monomer and carboxyl groups on the second. We will still form amide linkages, here with the different monomers alternating. (The polymer formed here is Kevlar®.) H N

O

NH O

Check Your Understanding The polymer whose structure is shown below is produced by a condensation reaction in which water is eliminated. Draw the structure of the corresponding monomer. H

[N

O

[

Copolymers Nylon and Dacron are two examples of copolymers made from more than one type of monomer. In these cases, because the monomers are arranged in a regular, alternating series, they can be further classified as alternating copolymers. There are other important types of copolymers as well (Figure 8.25). One of the more creative ways to engineer materials is to design block copolymers. A block copolymer has regions in the material where a single monomer unit is repeated, interspersed with other regions where a different monomer is the repeating unit. Spandex is an example of a block copolymer. In this material, some regions of the polymer are relatively stiff, whereas others are flexible. The resulting material is both strong and flexible, leading to widespread use in sports gear and fashion. Another engineering material that is a copolymer is commonly called ABS because the monomer units involved are acrylonitrile, butadiene, and styrene. This material is a Except for the alternating copolymer, these diagrams are schematic. The actual segments of each type of monomer would often be much longer than what is shown here.

Alternating copolymer

Block copolymer

Random copolymer

Graft copolymer

Figure 8.25 ❚ Copolymers can be grown with different arrangements of monomers. Here, the different colored circles represent two different monomers. For the most part, the names of these types of copolymers are self-explanatory. The different monomers alternate in an alternating copolymer or are randomly combined in a random copolymer. In a block copolymer, each type of monomer tends to be clumped together. Graft copolymers involve side chains of one polymer attached to a backbone of a different polymer. 270


graft copolymer in which both butadiene and acrylonitrile are attached to a polystyrene backbone. The resulting polymer offers a combination of the properties of its three constituents. Polystyrene provides ease of processing, glossiness, and rigidity. Acrylonitrile adds chemical resistance and hardness, and butadiene provides impact resistance. The specific composition can be tailored to a particular application. ABS is the standard material of choice for the strong plastic cases of computers, televisions, and other household electronic goods and sometimes is used as an alternative to PVC for pipes.

Physical Properties There are two key reasons for the importance of polymers as engineering materials. First and most obviously, polymers offer physical properties such as strength and elasticity that can be desirable in a wide range of uses. Second, those properties can be controlled or tailored to a greater degree than is usually possible in metals or other classes of materials. Through the choice of a particular polymer and the details of its synthesis and processing, materials engineers can choose the properties they need with a striking degree of specificity. In this section, we will look at a few examples of the types of choices that can be made. We begin by considering the thermal properties of polymers. In choosing a polymer for a specific design, you would obviously need to make sure that your finished object will be suitable for the temperature range in which it will be used. But an experienced engineer will also recognize the need to think about the way the object will be manufactured. Many plastic components are formed at high temperatures, and different polymers respond in dramatically different ways to heat. Polymers are often divided into two broad categories: thermoplastic and thermosetting. Thermoplastic polymers melt or deform on heating. This may seem like a weakness because it means that they are not suitable for high temperature applications. But a great many plastic objects, including children’s toys and bottles of many sorts, are generally used at ambient temperatures. So, the fact that they would melt if heated appreciably is not a major drawback. When we think about how to make these objects, which can come in rather complicated shapes, the ability to melt the polymer at reasonable temperatures becomes a major advantage. Depending on their complexity, products made from thermoplastic polymers are typically extruded or formed in molds or presses. The fact that the material softens or melts when heated allows shaping it into the desired form. Once cooled, the polymer solidifies and regains its structural properties. If we are trying to design an object that might need to be used at higher temperatures, though, the fact that thermoplastics soften and melt will make such materials a poor choice. Instead we might turn to thermosetting polymers, which can maintain their shape and strength when heated. The name “thermosetting” comes from the fact that these polymers must be heated to set or “lock in” their structures. But once this has been done, the materials offer increased strength and do not lose their shape upon further heating. Rather than being extruded, most thermosetting polymers are molded. The molecular origin of the difference between thermoplastic and thermosetting polymers is shown in Figure 8.26. The initial heating and setting of a thermosetting polymer produces a number of links between sites on the

In a thermoplastic polymer, chains interact only through intermolecular forces.

In thermosetting polymers, chains are tied together by actual chemical bonds.

There are levels of polymer structure, such as the formation of ordered regions with the polymer folded back onto itself, which we are not considering here.

Figure 8.26 ❚ The different properties of thermoplastic and thermosetting polymers result from the ways in which the polymer chains interact with one another.

8.6

Polymers

271

carbon backbone of different molecular chains. These links are referred to as crosslinks because they cross between and link individual molecular strands of the polymer. Chemically, these cross-links are additional covalent bonds that join the polymer chains to one another. Like most covalent bonds, they are strong enough that they do not readily fail upon heating. So the cross-linked polymer keeps its shape. An important example of the engineering importance of cross-linking in American industrial history is the discovery of vulcanization. In vulcanization, natural rubber is heated in the presence of sulfur. This produces cross-linking and leads to a harder material that is markedly more resistant to heat. Until vulcanization was discovered, natural rubber was difficult to use in applications such as automobile tires because it would become sticky when heated. Natural rubber is no longer widely used, having been replaced by synthetic forms. But the large-scale development of vulcanized tires and the design freedoms they afforded automotive engineers was an important component of the growth of the U.S. automobile and tire industries. Another critical physical property of polymers is elasticity. The ability of many polymeric materials to be stressed and deformed but return to their original shape is a characteristic that is often valuable to engineers. Fibers, in particular, must be elastic. Polymers that are particularly flexible and elastic are sometimes referred to as elastomers. Looking at the molecular structure of elastomers, we find that they tend to share some common traits. In particular, the geometry of the carbon backbone is often such that the elastomer forms an amorphous solid rather than a crystalline one. The forces between polymer molecules in an amorphous solid are not as strong as those in a more crystalline system, so they can be deformed and restored with less force.

Polymers and Additives Despite their wide range of useful properties, polymers alone often do not have the characteristics needed for a particular design. Fortunately, additives can be included in the material to control polymer properties further. Some additives have only modest mechanical effects. Pigments, for example, can be added solely to change the color of the material and don’t influence the underlying molecular structure of a polymer. Other additives play critical roles in the performance of the materials. PVC alone, for example, is too brittle for many applications. To rectify this, relatively small molecules called plasticizers are added to improve its flexibility. A plasticizer has to have several critical features. It must be capable of being incorporated into the solid polymer, so its structure should resemble the polymer being used. It must be nonvolatile, or it will escape from the solid too quickly and no longer impart the desired flexibility. Plasticizers do escape slowly, however, and often the smell associated with plastic materials arises from the release of trace amounts of these additives. Other additives that find regular use include antistatic agents, fillers, fire retardants, and light and heat stabilizers. There are many more features of polymers that you may learn about as you proceed in your engineering curriculum. From the perspective of chemistry, they serve as a tremendous example of the impressive structural variability possible with carbonbased molecules. Although other elements can form polymeric materials, it is the chemical characteristics of carbon that make the modern plastics industry possible.

INSIGHT INTO

8.7

The Invention of New Materials

Few components of research in the chemical sciences have a more profound impact on the everyday life of humans than the discoveries of new materials. There are, of course, many ways that new materials are found, including good fortune or serendipity.

272


The discovery of buckminsterfullerene, for example, shows several of the components of the way research in chemistry proceeds. The researchers at Rice University who discovered C60, led by Richard Smalley and Robert Curl, had already established the ability to study other elemental clusters, or small clumps of atoms. In doing so, they constructed sophisticated and unique instruments to generate small, gas-phase clusters as well as to detect their existence and measure the number of atoms in each cluster. When Harry Kroto visited the Smalley laboratories and suggested that they study carbon, there was no expectation that they would soon discover an entirely new field. Then they noticed that the experimental peak associated with a 60-atom carbon cluster, or C60, was unusually large. They decided to optimize the production of this species and attempted to explain its apparent unusual stability. Once they announced the discovery, other scientists repeated and extended the studies, and 5 years later, physicists at the University of Arizona discovered a means to make macroscopic quantities of C60 and other fullerenes. With this discovery, the world of nanotechnology was opened up and the road to potentially amazing materials was available for further research. The fact that C60 molecules are essentially tiny balls, for example, suggests that they might be useful as molecular scale bearings. Other materials have been discovered not by accident but by the deliberate application of well-established techniques. Traditional synthetic methods for new materials involve mixing solid precursors in ratios that are predicted to provide a desired stoichiometry. The mixture is then heated to allow the atoms in the separate solids to diffuse into each other. This process can be quite slow, so heating times are often long and temperatures are sometimes quite high. Often, the process has to be repeated several times to achieve the synthesis. Normally, the intermediate products must be ground into small particles, and hence this process has sometimes been dubbed the “heat-and-beat” method of synthesis. Sometimes, even established techniques such as heat-and-beat lead to unexpected results. Superconductivity arises in some materials when all resistance to the flow of electricity vanishes. Because such resistance represents a significant loss in all designs that use electricity, superconductors are very attractive materials. Traditional superconductive materials required extremely low temperatures. But in 1987, Paul Chu and his research team at the University of Houston discovered a new class of superconductors that lost all resistance at much higher temperatures. Here “much higher” is a relative term because the temperatures needed are still far below room temperature. The production of materials that are superconducting at room temperature remains a major goal of ongoing research in materials science. Such materials would lead directly to huge gains in efficiency in our use of electricity. The initial discovery of a group of materials such as the fullerenes most often occurs in the context of “pure” chemistry, carried out by researchers who are not necessarily striving toward specific practical applications. But identifying and developing practical uses for scientific knowledge is at the heart of engineering, of course. So it is quite common to see engineers become interested in newly discovered materials. A good example of this can be seen in the development of electrically conductive polymers, a field that currently holds great promise in many emerging applications. By now you should closely associate the word polymer with plastics. When you think of plastics, you expect them to be electrical insulators. After all, many plastics are routinely used for the insulation on wires. So how might we make a polymer into an electrical conductor? We learned in this chapter that electrical conductivity in metals and semiconductors is closely linked to the idea of electron mobility or delocalization: when electrons are free to move throughout a material, it can conduct electricity. That turns out to be the key to achieving conductivity in polymers, too. When we explored chemical bonding in Chapter 7, we invoked the concept of resonance in Lewis structures to allow drawing structures of molecules whose electrons were not specifically localized in a particular chemical bond. If we create a polymer for which we can draw extensive

The concept of careful doping that we considered in semiconductors can also be used in the design of superconducting materials.

8.7 The Invention of New Materials

273

resonance structures, then, there might be a possibility of achieving enough delocalization of electrons to lead to conductivity. Thinking back to Section 7.5, you may realize that one way to do this would be to set up an alternating sequence of double and single bonds down the backbone of the polymer. A prototypical example would be poly(acetylene), a portion of whose chain would have the structure shown below:

[

[

It should be apparent that a second resonance structure could be drawn in which the positions of all of the C!C and C"C bonds are reversed. This opens up the possibility of delocalizing significant numbers of electrons within a single polymer molecule—a big step toward conductivity. A few specific examples of electrically conductive organic materials have been known since the early 1900s. But they were generally regarded as curiosities until our understanding of both polymer structure and electrical conductivity evolved during the past century. That progress now makes it possible to explain earlier observations and also to begin to design new polymers rationally with specific electrical properties. The production of these conducting polymers involves a form of doping that is somewhat analogous to what we saw for inorganic semiconductors, and the fact that this doping can be carefully controlled holds great potential for tailoring the properties of individual polymers. In many ways, the possibilities are even greater for polymers. Only a handful of elements can be used as the underlying material for semiconductors, but the range of available monomers and polymers is much broader. Applications of conducting polymers are currently being pursued in several areas, most notably as organic light-emitting diodes (OLEDs). Different preparations can produce polymers that will display different colors. Because the polymers can be dissolved, they can be manipulated in intriguing ways. A polymer solution can be deposited through a process similar to ink-jet printing, and adding contacts can produce a highly customized display that can be applied to almost any surface. The development of conducting polymers has seen important contributions from both chemists and engineers, as is usually the case. The interplay between the discoveries of the scientist, driven largely by curiosity, and the advances of the engineer seeking a solution to a specific problem is synergistic. Progress on either side fuels the further thinking of both groups and spurs the next generation of advances.

FOCU S O N

PRO BLEM S O LVI NG

Question The metals in Groups 11 and 12 all have face-centered cubic structures. They are also among the least reactive metals. Aside from looking up the density, what could you look up to determine which metal was the densest, and how would that information help you to determine the density?

Strategy This question has both a conceptual component and a computational component. First, we need to understand how a lattice is related to density, and then we need to identify the computational pathway that relates structures at the atomic level to densities at the macroscopic level.

Solution Knowing that the structure is face-centered cubic tells us about the geometry of the atoms—how they are connected. If we can look up the length of the cube edge, we can calculate the number of atoms in a given volume (related to the volume of the unit cell). Then we would need to look up the atomic mass to convert between number of atoms in the given volume and the mass of those atoms. Density is mass per unit of volume, so these steps would provide a density estimate that could be used to estimate which metal would have the greatest density.

274


S U M M ARY Understanding chemical bonding is vitally important in being able to devise models for materials, including those important in engineering applications. In most materials, however, the forces that must be considered extend beyond a localized bond between two atoms. Thus an understanding of the interactions that give rise to the bulk properties of materials requires us to consider additional ideas. For solid materials, the packing of the constituent particles into a regular array, or lattice, often provides the first simplifying component for building a model. There are important disordered, or amorphous, solids, but the order in crystalline solids makes them much easier to understand. A crystalline solid can be described at the atomic or molecular level in terms of its packing efficiency or coordination number. To understand many properties of solids, including electrical conductivity, we must consider the packing of the atoms and also the way in which the electrons are shared or distributed. Thus we are able to extend our notion of chemical bonding and consider metals and other extended systems, using models with varying levels of sophistication. The simple electron-sea model accounts for some metallic behaviors, but a more complete bonding model called band theory allows an additional level of detail. Using band

theory, we can also understand the differences between metals, semiconductors, and insulators as well. The ability to manipulate materials—particularly semiconductors—by doping them with small levels of impurities to obtain desired characteristics is generally predicated on band theory. Crystalline solids are not the only form of condensed phase, however, and to consider other solids or liquids, we must define the nature of the interaction between molecules. Several types of intermolecular forces have been discerned, including dispersion, dipole–dipole forces, and hydrogen bonding. The strength of these intermolecular forces helps explain observable macroscopic properties such as vapor pressure, boiling point, and surface tension. From an engineering perspective, one class of materials for which the nature of intermolecular interactions plays an enormous role is polymers. Here, we must consider the chemical bonding within a single polymer molecule and also the interaction between polymer chains via intermolecular forces. Careful consideration of these factors can enable the production of materials with specific physical properties. Thus when a materials engineer needs a plastic for a specific application, the choice of monomers and how they are polymerized provide the two critical variables that can be manipulated to lead to a polymer with the required physical properties.

K EY TE RM S acceptor level (8.3)

crystalline solid (8.2)

n-type semiconductor (8.3)

addition polymerization (8.6)

degree of polymerization (8.6)

packing efficiency (8.2)

adhesion (8.5)

dipole–dipole forces (8.4)

phase diagram (8.1)

alternating copolymer (8.6)

dispersion forces (8.4)

p-n junction (8.3)

amorphous (8.2)

donor level (8.3)

polarizability (8.4)

antibonding molecular orbital (8.3)

doping (8.3)

p-type semiconductor (8.3)

atactic polymer (8.6)

ductile (8.3)

sea of electrons model (8.3)

band gap (8.3)

dynamic equilibrium (8.5)

simple cubic (8.2)

band theory (8.3)

face-centered cubic (8.2)

superconductivity (8.7)

block copolymer (8.6)

graft copolymer (8.6)

surface tension (8.5)

body-centered cubic (8.2)

hydrogen bonding (8.4)

syndiotactic polymer (8.6)

bonding molecular orbital (8.3)

insulator (8.3)

thermoplastic polymer (8.6)

cohesion (8.5)

intermolecular forces (8.4)

thermosetting polymer (8.6)

condensation polymer (8.6)

isotactic polymer (8.6)

unit cell (8.2)

conduction band (8.3)

malleable (8.3)

valence band (8.3)

coordination number (8.2)

nanotube (8.1)

vapor pressure (8.5)

copolymer (8.6)

normal boiling point (8.5)

volatile (8.5)



INSIGHT INTO Carbon 8.1 How many solid forms of elemental carbon are known? 8.2 Why is the C60 form of carbon called buckminsterfullerene?

■


8.3 What property of diamond leads to the most engineering applications? Which types of applications would benefit from this property? 8.4 Which variables must be controlled and in what manner to produce diamonds from graphite?


275

8.5 What is the relationship between the structures of buckminsterfullerene and carbon nanotubes?

8.25 How does the sea of electrons model of metallic bonding explain why metals are good conductors of electricity?

8.6 Use the web to look up information on nanotubes. Distinguish between single-walled and double-walled nanotubes.

8.26 What is the key difference between metallic bonding (in the sea of electrons model) and ionic bonding (as described in Chapter 7) that explains why metals conduct electricity and ionic solids do not?

8.7 Use the web to look up the experimental conditions required to synthesize buckminsterfullerene in an electric arc. Would running an arc in air work? Condensed Phases—Solids 8.8 What is the difference between an amorphous solid and a crystalline solid? 8.9 Define packing efficiency.

8.27 What is the difference between a bonding orbital and an antibonding orbital? 8.28 Describe how the combination of atomic orbitals gives rise to bands in the limit of large numbers of atoms. 8.29 Draw a depiction of the band structure of a metal. Label the valence band and conduction band.

8.10 Using circles, draw regular two-dimensional arrangements that demonstrate low packing efficiency and high packing efficiency.

8.30

8.11 Using pentagons, draw arrangements that demonstrate low packing efficiency and higher packing efficiency.

8.31 The conductivity of semiconductors increases as temperature is increased. Using band theory, explain this observation.

8.12 Using circles, show that a cubic structure (squares in 2-D) has a lower coordination number than a hexagonal structure. 8.13

■ Calculate

the packing efficiencies for both two-dimensional arrangements shown in Figure 8.3.

8.14 Explain why hcp and ccp structures have the same coordination number. 8.15 What is the coordination number of atoms in the diamond structure? 8.16 Explain why graphite is described as having a layered structure. 8.17

8.18

8.19

8.20

Polonium is the only metal that forms a simple cubic crystal structure. Use the fact that the density of polonium is 9.32 g/cm3 to calculate its atomic radius. ■

Iridium forms a face-centered cubic lattice, and an iridium atom is 271.4 pm in diameter. Calculate the density of iridium.

Manganese has a body-centered cubic unit cell and has a density of 7.88 g/cm3. From this information, determine the length of the edge of the cubic cell. ■

8.21 The spacing in a crystal lattice can be measured very accurately by X-ray diffraction, and this provides one way to determine Avogadro’s number. One form of iron has a bodycentered cubic lattice, and each side of the unit cell is 286.65 pm long. The density of this crystal at 25ºC is 7.874 g/cm3. Use these data to determine Avogadro’s number. Metals and Metallic Bonding 8.22 How many electrons per atom are delocalized in the sea of electrons model for the following metals? (a) iron, (b) vanadium, (c) silver 8.23 One important physical property of metals is their malleability. How does the sea of electrons model account for this property? 8.24 The sea of electrons model is not generally used for quantitative predictions of properties. What factors are left out of this model that might prevent quantitative precision?

276

8.32 Use the web to find out what a “III–V” semiconductor is. How is this type of semiconductor related to silicon? 8.33 Explain how doping affects the electronic properties of a semiconductor. 8.34 What type of atom is needed as a dopant in an n-type semiconductor? Why is it called n-type? 8.35 What type of atom is needed as a dopant in a p-type semiconductor? Why is it called p-type? 8.36 Is an n-type semiconductor actually negatively charged? 8.37 How does a p-n junction serve as a voltage-gated switch? 8.38

■

■ Europium forms a body-centered cubic unit cell and has a density of 4.68 g/cm3. From this information, determine the length of the edge of the cubic cell.


■ In terms of band theory, what is the difference between a conductor and an insulator? Between a conductor and a semiconductor?

Suppose that a device is using a 15.0-mg sample of silicon that is doped with 1 × 10–5% (by mass) phosphorus. How many phosphorus atoms are in the sample? ■

Intermolecular Forces 8.39 What is an instantaneous dipole? 8.40 Why are dispersion forces attractive? 8.41 If a molecule is not very polarizable, how will it respond to an external electric field? 8.42 What is the relationship between polarizability and dispersion forces? 8.43 Most gaseous compounds consist of small molecules, while polymers are never gaseous at room temperature. Explain this observation based on intermolecular forces. 8.44 Why are dipole–dipole forces typically stronger than dispersion forces? 8.45 Under what circumstances are ion–dipole forces important? 8.46

Which of the following compounds would be expected to form intermolecular hydrogen bonds in the liquid state? (a) CH3OCH3 (dimethyl ether), (b) CH4, (c) HF, (d) CH3CO2H (acetic acid), (e) Br2, (f ) CH3OH (methanol) ■

8.47 What is the unique feature of N, O, and F that causes them to play a role in hydrogen bonding? 8.48

What type of intermolecular forces must be overcome in converting each of the following from a liquid to a gas? (a) CO2, (b) NH3, (c) CHCl3, (d) CCl4 ■

■


8.49

8.50

■ Identify the kinds of intermolecular forces (London dispersion, dipole–dipole, or hydrogen bonding) that are the most important in each of the following substances: (a) methane (CH4), (b) methanol (CH3OH), (c) chloroform (CHCl 3 ), (d) benzene (C 6 H 6 ), (e) ammonia (NH 3 ), (f ) sulfur dioxide (SO2)

Rank the following in order of increasing strength of intermolecular forces in the pure substances. Which exists as a gas at 25°C and 1 atm? (a) CH3CH2CH2CH3 (butane), (b) CH3OH (methanol), (c) He ■

8.51 Carbon tetrachloride (CCl4) is a liquid at room temperature and pressure, whereas ammonia (NH3) is a gas. How can these observations be rationalized in terms of intermolecular forces?

8.63 Draw the meniscus of a fluid in a container where the interactions among the liquid molecules are weaker than those between the liquid and the container molecules. 8.64 When water is in a tube such as a buret, in which direction does the meniscus curve? What does this observation say about the relative magnitude of adhesion and cohesion in that system? Polymers 8.65 If shown structures of monomers that might polymerize, what would you look for to determine if the type of polymerization would be addition? 8.66 Why is temperature increased to start most addition polymerization reactions?

8.52 Explain from a molecular perspective why graphite has properties that are useful for lubrication.

8.67 What is meant by the term chain termination in a polymerization reaction?


8.68 The monomer of polyvinylchloride is H2C"CHCl. Draw an example of an isotactic PVC polymer.

8.53 Describe how interactions between molecules affect the vapor pressure of a liquid. 8.54 What makes a chemical compound volatile? 8.55

Answer each of the following questions with increases, decreases, or does not change.

■

(a) If the intermolecular forces in a liquid increase, the normal boiling point of the liquid . (b) If the intermolecular forces in a liquid decrease, the vapor pressure of the liquid . (c) If the surface area of a liquid decreases, the vapor pressure . (d) If the temperature of a liquid increases, the equilibrium vapor pressure . 8.56 Why must the vapor pressure of a substance be measured only after dynamic equilibrium is established? 8.57

Which member of each of the following pairs of compounds has the higher boiling point? (a) O2 or N2, (b) SO2 or CO2, (c) HF or HI, (d) SiH4 or GeH4

8.58

■ Predict the order of increasing vapor pressure at a given temperature for the following compounds: (a) FCH2CH2F, (b) HOCH2CH2OH, (c) FCH2CH2OH

■

8.59 A substance is observed to have a high surface tension. What predictions can you make about its vapor pressure and boiling point? 8.60

8.61

■ Suppose that three unknown pure substances are liquids at room temperature. You make vapor pressure measurements and find that substance Q has a pressure of 110 torr, substance R has a pressure of 42 torr, and substance S has a pressure of 330 torr. If you slowly increase the temperature, which substance will boil first and which will boil last?

Suppose that three unknown pure substances are liquids at room temperature. You determine that the boiling point of substance A is 53ºC, that of substance B is 117ºC, and that of substance C is 77ºC. Based on this information, rank the three substances in order of their vapor pressures at room temperature.

■

8.62 Rank the following hydrocarbons in order of increasing vapor pressure: C2H6, C10H22, CH4, C7H16, C22H46.

■


8.69 Why is there no isotactic or syndiotactic form of polyethylene? 8.70 Why are isotactic or syndiotactic polymers often more attractive for materials development? 8.71 What are the products of a condensation polymerization reaction? 8.72

■ Nylon-6

is made from a single monomer: O

H2N

OH

Sketch a section of the polymer chain in Nylon-6. 8.73 Distinguish between a block copolymer and a graft copolymer. 8.74 Use the web to look up the polymers used for six synthetic fibers. Classify the polymers as (a) addition or condensation polymers and (b) as alternating, block or graft copolymers if they are copolymers. 8.75 What is the key physical property that characterizes a thermoplastic polymer? 8.76 What conditions would prohibit the use of a thermoplastic polymer? 8.77 What happens molecularly in a thermosetting polymer when it is heated? 8.78 What structural characteristics are needed for additives such as plasticizers? 8.79 Use the web to find out what pencil erasers are made of. Would it be possible to form a new shape of an eraser by melting it and pouring it into a mold? 8.80 Use the web to find an application where the presence of isotactic, syndiotactic, or atactic polymers affects the physical properties important for an application of a polymer. INSIGHT INTO the Invention of New Materials 8.81 Use the web to investigate the story of the discovery of buckminsterfullerene. What was the initial reaction to the announcement of this material?


277

grill on a hot summer day when the temperature is 95ºF. (c) What implications might your answer to (b) have for an engineer designing propane storage tanks?

8.82 What was the observation that caused Smalley and coworkers to investigate further and ultimately discover C60? 8.83 In the “heat-and-beat” method of materials synthesis, why is heat needed? 8.84 Many experts feel that applications of OLED technology will soon become common. Do a web search to identify emerging uses for OLEDs. Can you find any available commercial products using OLED displays?

T (°C)

–100

P (torr)

22.4

–80

–60

–40

–20

0

20

100.6 328.8 856.9 1888.6 3665.6 6445.9

Additional Problems

8.90 Geckos are known for their ability to climb smooth vertical walls or windows. Recent research has shown that this unusual trait results from van der Waals forces between the surface and the tips of tiny hairlike setae in the Gecko’s feet. Many materials science and engineering groups are working to develop tapes and other adhesives designed to mimic this behavior. Do a web search to find information about such efforts and write a short summary explaining how one resulting adhesive is related to ideas we explored in this chapter.

8.87

8.91

8.85 Use the web to find information about new materials being considered in potential battery technology for the future. Describe the characteristics that are described as helpful for applications in batteries. ■ Which of the following monomers might you expect would lead to a conducting polymer? (a) H!C#C!CH3, (b) H2C"CH!CH3

Use the vapor pressure curves illustrated here to answer the questions that follow. ■

Vapor pressure (mm Hg)

8.86

900 800 700 600 500 400 300 200 100 0

Carbon disulfide

Ethanol Heptane

10

20

30

40

50

60

70

80

90 100 110

Temperature (°C)

(a) What is the vapor pressure of ethanol (C2H5OH) at 60°C? (b) Considering only carbon disulfide (CS2) and ethanol, which has the stronger intermolecular forces in the liquid state? (c) At what temperature does heptane (C 7H 16) have a vapor pressure of 500 mm Hg? (d) What are the approximate normal boiling points of each of the three substances? (e) At a pressure of 400 mm Hg and a temperature of 70°C, is each substance a liquid, a gas, or a mixture of liquid and gas? 8.88 Cylinders of compressed gas are equipped with pressure gauges that allow the user to monitor the amount of gas remaining. But such gauges are not useful for substances like propane or carbon dioxide, which are stored as liquids. So tanks of propane or carbon dioxide are sometimes mounted on scales to allow the user to see the rate of consumption. Use ideas from this chapter to explain why this is so. 8.89

278

■ The following data show the vapor pressure of liquid propane as a function of temperature. (a) Plot a vapor pressure curve for propane and use it to estimate the normal boiling point. (b) Use your curve to estimate the pressure (in atm) in the propane tank supplying fuel for a gas barbecue


If you place 1.0 L of ethanol (C2H5OH) in a room that is 3.0 m long, 2.5 m wide, and 2.5 m high, will all the alcohol evaporate? If some liquid remains, how much will there be? The vapor pressure of ethanol at 25°C is 59 mm Hg, and the density of the liquid at this temperature is 0.785 g/cm3.

■

8.92 An environmental engineering team is brought in to investigate a small oil spill caused by a leak in a crude oil pipeline. They find that the average molar mass of a sample collected from the spilled oil is notably higher than that of the oil in the pipeline. Nonetheless, they are able to verify that the oil in the spill originated in the pipeline. Use ideas related to intermolecular forces to explain the difference between the spill oil and the pipeline oil. 8.93 Do a web search to find out how p- and n-type semiconductors can be used in designing solar panels. Write a short paragraph explaining what you learn. What new types of materials are being considered for use in solar panel systems? 8.94 Hydraulic fluids are liquids moving in a confined space under pressure, and they can serve many purposes in engineering designs. One use of hydraulic fluids is lubrication. Suppose you need to choose a hydrocarbon for use as a lubricant. What size molecules would provide high viscosity for a lubricant in a design? 8.95 It is possible to make syndiotactic polystyrene, although most polystyrene is atactic. Syndiotactic polystyrene melts at 270°C while atactic polystyrene melts closer to 240°C. (a) Why might this difference in structure result in different boiling points? (b) If your supplier charges 20% more for syndiotactic polystyrene, what factors of a design might call for using this material despite the higher cost? 8.96 A business manager wants to provide a wider range of p- and n-type semiconductors as a strategy to enhance sales. You are the lead materials engineer assigned to communicate with this manager. How would you explain why there are more ways to build a p-type semiconductor from silicon than there are ways to build an n-type semiconductor from silicon? 8.97 The doping of semiconductors can be done with enough precision to tune the size of the band gap in the material. Generally, in order to have a larger band gap, the dopant

■


should be smaller than the main material. If you are a materials engineer and need a semiconductor that has lower conductivity than pure silicon, what element or elements could you use as your dopant? (You do not want either an n- or a p-type material.) Explain your reasoning.

FOCUS ON PROBLEM SOLVING EXERCISES

high. What types of measurements must be made in the laboratory to determine whether the degree of polymerization is acceptable? Cumulative Problems 8.102 How is polarizability related to the periodic trends of elements in the periodic table?

8.98

If you know the density of a material and the length of the edge of its cubic lattice, how would you determine if it is face-centered cubic, body-centered cubic, or simple cubic? Would you have to look up any information?

8.103 In previous chapters, we have noted that cryolite is used in refining aluminum. Use the web to look up what this addition does for the process, and relate it to the concepts of intermolecular forces from this chapter.

8.99

If a pin is run across a magnet a number of times, it will become magnetized. How could this phenomenon be used as an analogy to describe polarization in atoms and molecules?

8.104 Use the web to look up the percentage of dopant for a commercially available p-type semiconductor. Imagine that you were setting up a process for doping 1 metric ton of silicon with this dopant. (a) What mass would be required? (b) What would be the mole fraction of the dopant?

8.100 You go into the laboratory to look for a squirt bottle containing acetone. You find two unlabeled bottles with different colored tops suggesting they are different liquids. Unfortunately, you have a terrible cold and cannot tell by smell which one might be acetone. What simple test could you use to determine which liquid is acetone and which is water? How does this test tell you this information? 8.101 A materials engineer, with an eye toward cost, wants to obtain a material whose degree of polymerization is

■


8.105 Use the web to look up the percentage of dopant for a commercially available n-type semiconductor. Imagine that you were setting up a process for doping 1 metric ton of silicon with this dopant. (a) What mass would be required? (b) What would be the mole fraction of the dopant?


279

9

Energy and Chemistry

OUTLINE 9.1

INSIGHT INTO Energy Use

and the World Economy 9.2

Defining Energy

9.3

Energy Transformation and Conservation of Energy

9.4

Heat Capacity and Calorimetry

9.5

Enthalpy

9.6

Hess’s Law and Heats of Reaction

9.7

Energy and Stoichiometry

9.8

INSIGHT INTO Batteries

Wind farms like the one shown here are one of several alternative energy technologies that have grown in importance in recent years. Despite these advances, fossil fuels will continue to provide the bulk of our energy for the foreseeable future. Kevan O'Meara, 2009/Shutterstock.com

I

f you list applications of chemistry that are important in your life, many of them would be likely to involve energy and energy transformation. In Chapter 4, we looked at the burning of fuels as an application of chemical combustion. All the batteries you depend on to power portable electrical devices are based on chemical reactions as well. In this chapter, we will take a closer look at energy and its role in chemistry. We begin by focusing on the production and use of energy in North America, a component of life in the developed world that we’ll refer to as the energy economy.

Chapter Objectives


280

After mastering this chapter, you should be able to ❚ explain the economic importance of conversions between different forms of energy and the inevitability of losses in this process. ❚ define work and heat using the standard sign conventions. ❚ define state functions and explain their importance. ❚ state the first law of thermodynamics in words and as an equation. ❚ use calorimetric data to obtain values for DE and DH for chemical reactions. ❚ define DHf ° and write formation reactions for compounds.

❚ explain Hess’s law in your own words. ❚ calculate DH ° for chemical reactions from tabulated data. ❚ describe some important design considerations in choosing a battery for a specific application.

INSIGHT INTO

9.1

Energy Use and the World Economy

25,000

Energy use per dollar of GDP

Canada 20,000 South Korea 15,000 U.S.

Taiwan

Mexico

10,000

5,000

0 1980

1985

1990

1995

2000

2005

Data are from the Energy Information Administration, http://www.eia.doe.gov.

Few factors reveal the wealth of a country or region more vividly than its energy consumption. The central role of the United States in the world economy is reflected in the fact that the nation accounts for 24% of world energy use. Canada, Japan, and Western Europe are also large users of energy. Energy use is closely connected to a broad array of societal, economic, and political factors. For example, in the early 1990s, Pacific Rim countries such as Thailand and South Korea had growing economies and substantial increases in their energy use. Yet during this time, energy use worldwide was stagnant, as shrinking economies in Eastern Europe offset the increases elsewhere in the world. Figure 9.1 illustrates the link between economic growth (as measured by gross domestic product) and energy use. Data are plotted for several countries as energy consumption per dollar of GDP from 1980 to 2006. It is clear from the figure that the lines shown are all nearly horizontal. This confirms the correlation between economic performance and energy consumption because a horizontal line in this plot implies that energy use is proportional to GDP. But the figure also shows that some countries are more energy-dependent than others because the various lines are not identical. This is partly due to differences in the efficiency with which different countries use energy and partly due to climate and other factors. We can learn more about the pivotal role of energy in an economic system by examining the sources and uses of energy. Although there are differences between countries, we’ll look at the United States as an example of energy production and use in an industrialized nation. Figure 9.2 is a visual representation of the pattern of energy production and use in the United States in 2007. At first glance, this figure is somewhat complicated, but it can be analyzed to reveal a great deal of information. Start by looking at each end: the left side represents sources of energy, and the right side shows consumption.

GDP is usually defined as the total value of all goods and services produced annually in a region.

Figure 9.1 ❚ The graph demonstrates the link between energy consumption and gross domestic product (GDP), which is a common indicator of economic strength. The data are plotted in units of Btu per U.S. dollar, with all data converted to constant (2000) dollar values. The fact that these lines are all essentially horizontal confirms the connection between energy use and economic strength.

Year

9.1 Energy Use and the World Economy

281

Co

al 2

Nat

ural

Crud e NGP

gas

Exports 5.36 Other Exports 2.43

3.4

8 Fossil fuels 56.50 Domestic production 71.71

19.8

2

oil 1

0.80

L 2.4

available at http://www.eia.doe.gov/emeu/aer/.

Figure is from the Energy Information Administration’s Annual Energy Review. The most recent report is

Petroleum 2.93

0

Coal 22.77

Supply 106.96

Other Imports 5.90

Imports 34.60

Commercial 18.43

Natural gas 23.64 Fossil fuels Petroleum 39.82 86.25 Consumption 101.60

Nuclear electric power 8.41 Renewable energy 6.80 Petroleum 28.70

.75

tial 21

en Resid

Nuclear electric power 8.41 Renewable energy 6.83

Industrial 32.32

Transpo r

tation 2

9.10

Stock Change and Other 0.65

Figure 9.2 ❚ Energy production and consumption (in quadrillion Btu) in the United States during the year 2007. The discussion in the text explains how to read this complex figure, which contains an enormous amount of information about the energy economy. Data come from the Department of Energy and do not always add up exactly as expected due to rounding and other issues.

Some assumptions are built into the division of uses into these four categories, so absolute numbers are not as important as comparative ones.

In economic terms, energy is a commodity. So when supplies are low or demand is high, prices tend to increase.

282

Thus the whole figure represents the flow of resources through the U.S. energy economy. All numbers on the figure are given in units of quadrillion Btu. (A quadrillion is 1015, and the Btu is a unit of energy. The letters stand for British thermal unit, and 1 Btu =1054.35 J.) Starting in the upper left, we can examine the sources of energy production. Coal accounted for 23.48 quadrillion Btu, or 32.7% of the total domestic production of 71.71 quadrillion Btu. Natural gas and crude oil provided 27.6% and 15.1% of the total, respectively. Smaller contributions came from nuclear energy and renewable energy sources, which include hydroelectric power, wood, solar energy, and wind. Energy imports appear at the bottom left of the figure, where we see that the main imports are crude oil and related petroleum products. Combining the domestic production and the imports gives the total supply of energy, which was 106.96 quadrillion Btu for 2007. Moving to the right, the figure divides this supply to show how energy was consumed. The four main components of use are classified as (1) residential, (2) commercial, (3) industrial, and (4) transportation, accounting for roughly 21, 18, 32, and 29% of total energy use, respectively. Nearly half of all domestic energy use goes into producing electricity. The generation and consumption of electricity is detailed in Figure 9.3. (The form of this diagram is similar to that of Figure 9.2, so you should examine it in the same way.) This figure shows that conversion losses account for nearly two-thirds of the energy consumed to generate electricity. These conversion losses in electricity generation mean that more than 25% of the total domestic energy consumption is lost from the economy. This raises some important questions: How is energy from other sources converted to electricity, and why is this conversion necessary? Why is so much energy lost or wasted in this process, and how can we minimize or eliminate this loss? Other questions arise when considering the relative roles of various fossil fuels in the energy economy. Figure 9.4 shows the history of U.S. domestic consumption of energy from various sources. It’s clear that our consumption of fossil fuels has grown steadily throughout our country’s entire history. But this growth can be sporadic, and there have been periods in which consumption of petroleum and natural gas has actually decreased substantially. A variety of factors contribute to these trends

Chapter 9 Energy and Chemistry

Fossil fuels 29.59

Other gases 0.17

Conversion losses 27.15 Energy consumed to generate electricity 42.09

Natural gas 7.72

Plant use 0.75 T&D losses 1.34

Petroleum 0.72

Residential 4.7

Gross Net generation generation of electricity of electricity 14.94 14.19

Nuclear electric power 8.41 Renewable energy 3.92

5

Retail sales Commercia l 4.58 12.79

End use 13.28

Industr

ial 3.4 3

Electricity imports 0.11 Unaccounted for 0.32

Other 0.17

Electricity exports Direct Transuse portation 0.08 0.49 0.03

Data from the Energy Information Administration’s Annual Energy Review.

Coal 20.99

Figure 9.3 ❚ Visual summary of the generation and consumption of electricity in the United States during the year 2007, presented in a manner similar to Figure 9.2.

in consumption, notably the availability of raw material and the price and availability of imported fuels. The decrease in petroleum consumption in the early 1980s, for example, coincides with the recession of that era, a period in which decreased economic activity led to lower demand for oil. More recently, higher oil prices have driven interest in renewable energy sources. Because our interest here is chemistry, though, we will want to ask whether there are more fundamental differences among these

120.00

Biofuels Projections

History

Quintillion Btu

80.00

Liquids

60.00 Natural Gas 40.00 Nuclear 20.00 Coal 0.00 1980

1985

1990

1995

2000

2005 Year

2010

2015

2020

2025

2030

Data are from the Energy Information Administration's Annual Energy Outlook, 2009.

Renewables

100.00

Figure 9.4 ❚ U.S. consumption of energy from various sources changes slowly over time. The graph shows actual consumption for 1980–2007 and projected use for 2008–2030.

9.1 Energy Use and the World Economy

283

fuels in terms of energy production and transformation. What is the role of chemistry in deciding which of these fossil fuels is most attractive in a given set of circumstances? To address all of the questions raised here we must consider some aspects of thermochemistry—the study of the energetic consequences of chemistry.

9.2

Defining Energy

The questions raised in Section 9.1 lead us into a study of energy. Most people have an intuitive feel for the concept of energy, and we see the word every day. Breakfast cereals advertise their ability to provide energy in the morning. In sports, substitute players are described as boosting their team’s energy. The energy crisis brought on by oil embargoes in the 1970s is a critical component of the history of that decade, and energy remains central to important political issues in the 21st century. This widespread use of the concept of energy underscores its importance. It also suggests that we’ll need to be very careful in defining energy and related terms because it is apparent that in casual use the word “energy” conveys many related but subtly different meanings.

Forms of Energy Most of the energy we encounter can be placed into two broad categories: potential energy and kinetic energy. Potential energy is associated with the relative position of an object. For example, a roller coaster gains potential energy as it is pulled up the initial slope because it acquires a higher position relative to the ground and gravity is attracting it downward. But gravity is not the only force responsible for objects having energy associated with their positions. The attraction and repulsion between electrical charges also lead to potential energy, as we’ve seen in our consideration of the structure of atoms and molecules. Kinetic energy is associated with motion. When the roller coaster heads down the first hill, it transforms potential energy into kinetic energy. You probably are familiar with Equation 9.1, which provides a mathematical definition for the kinetic energy of an object in terms of its mass (m) and velocity (v). 1 mv2 Kinetic energy = — 2

Chemical energy is a form of potential energy.

284

(9.1)

The roller coaster example shows kinetic and potential energy on a macroscopic scale. But on a microscopic level, all substances and objects, from fuels to the paper on which these words are printed, have these same forms of energy. Any substance or object is composed of atoms and molecules. These atoms and molecules have kinetic energy associated with their constant motion, and they have potential energy due to the various forces they exert on one another. The combined kinetic and potential energies of the atoms and molecules that make up an object constitute its internal energy. Thus the roller coaster cars have three basic forms of energy: kinetic (from their motion), potential (from their position relative to the ground), and internal (from the molecules that compose the materials from which they are made). Much of the internal energy of an object, be it a roller coaster car or a piece of coal, is associated with the potential energy arising from the relative positions of the atoms that make up the object. As we saw in Chapter 7, the making or breaking of chemical bonds changes this potential energy. In most chemical reactions, bonds are broken in the reactants and new bonds are formed in the products. If the amount of energy liberated in bond making is greater than that consumed in bond breaking, the overall process releases energy. This energy release is often referred to as chemical energy, and the harnessing of chemical energy is an important aspect of the overlap between chemistry and engineering. Occasionally, we find it useful to have more specific terms for various forms of energy as well. Radiant energy is associated with light or electromagnetic radiation.


The radiant energy of the Sun is ultimately responsible for the majority of energy resources on this planet. Mechanical energy is associated with the movement of macroscopic objects. Thermal energy arises from the temperature of an object. (We can associate thermal energy with the molecular level motion of atoms and molecules.) Electrical energy results from moving charge—usually electrons in a metal. Nuclear energy, which can be released in nuclear fusion and fission processes, is a form of potential energy associated with the arrangement of protons and neutrons in atomic nuclei.

Heat and Work Although we can use a wide variety of classifications for types of energy, all energy flow is either heat or work. We’ll need to understand both forms of energy transfer to assess the energy economy of the world and the role of chemistry in that economy. Heat is the flow of energy between two objects, from the warmer one to the cooler one, because of a difference in their temperatures. Thus if we are speaking carefully, heat is a process and not a quantity. Although we routinely hear statements such as “turn up the heat,” heat is not an entity we can pump into a room or a cup of coffee. An object does not possess heat. In a strictly scientific sense, a furnace does not produce heat but rather a body of warm air or hot water that has a higher temperature than the cool air in a room. What emerges from the vent on the floor is not “heat,” but warm air. Although these distinctions are essentially semantic, they can be very important in many cases. Work is the second form of energy transfer. Work is the transfer of energy accomplished by a force moving a mass some distance against resistance. Lifting a set of roller coaster cars up a hill against the pull of gravity is an example of work. When we consider macroscopic examples, we are typically viewing work in terms of mechanical energy. Work, however, encompasses a wider range of phenomena than just mechanical movement of macroscopic objects. The most common type of work we will encounter in chemical processes is pressure-volume work (PV-work). When a gas expands, it can do work. If an inflated balloon is released before it is tied off, it flies around as the gas inside the balloon expands into the large volume of the room. Because the flying balloon has mass, it is easy to see that the expanding gas is doing work on the balloon: this is pressure-volume work. For a more productive example of work being done by a chemical reaction, we might look at the burning of gasoline in a car engine. We learned in Section 4.1 that gasoline is actually a complex mixture of hydrocarbons. The energy needed to propel a car is released by the combustion of those hydrocarbons in the engine cylinders. Hydrocarbon + O2(g) : CO2(g) + H2O(g) This combustion produces carbon dioxide and water vapor, and those gases do PVwork as they expand against the piston in the cylinder. This PV-work is then transmitted through the drive train to move the car.

Energy Units We use a wide variety of units when measuring and discussing energy, and various disciplines of science and engineering tend to favor particular choices of units. The SI unit of energy is the joule ( J), and 1 joule is equal to 1 kg m2/s2. Realizing that work should have units of energy helps us to make sense of this unit. Work is force times distance, and force is mass times acceleration. Thus, work is Mass × acceleration × distance or 2 m × m = kg m kg × —— ——— s2 s2

9.2 Defining Energy

285

It’s often handy to develop an intuition for the magnitude of a particular unit. A joule is roughly the amount of energy it takes to raise a 1-kg (about 2 lb) book 10 cm (4 in.) above a tabletop. If we think about chemistry on a molecular level, a joule is a very large amount of energy; breaking a single chemical bond typically requires only about 10–18 J. But we seldom break only one bond at a time. If we look at a reaction from the macroscopic viewpoint, we might consider breaking a mole of bonds. In that case, the joule seems like too small a measure of energy, and we might use kJ (or kJ/mol) instead. Breaking one mole of C!H bonds, for example, requires about 410 kJ. Many older units for energy relied on some readily observable property for their original definition. Thus the Btu, which is still widely used in several engineering disciplines, was first defined as the amount of energy needed to raise the temperature of 1 lb of water by 1°F. Another traditional unit for energy is the calorie, which was originally defined as the amount of energy required to heat 1 g of water from 14.5 to 15.5°C. The joule has become the most widely accepted unit for energy, and so these other units are now defined in terms of it. A calorie is defined as 4.184 J, and a Btu is defined as 1055 J. A source of frequent confusion in comparing energy units is that the Calorie reported for foods is actually a kilocalorie. (The food Calorie should be written with an uppercase “C.”) Thus, 1 Calorie is actually equivalent to 4184 J, or 4.184 kJ. The energy typically consumed by large economies is enormous, as we saw when analyzing Figures 9.2 and 9.3.

9.3

Careful definition of the system and its surroundings is also important in many engineering problems.

286


We have introduced several ways to categorize energy. But these multiple forms of energy are not all equally useful, so in many cases it is desirable to transform energy from one form into another. For example, the lighting in your room is provided by electricity, but that electricity was probably generated by the release of chemical energy through the combustion of coal. Unless you want to try to light your room by burning a chunk of coal, you need a way to harness the chemical energy released as the coal burns and then convert it to electrical energy. That electrical energy must then be conveyed to your room, where your light bulbs convert it into radiant energy. We have already noted the reality of energy waste in these processes in Section 9.1. Now we begin to consider the laws of nature that apply when one form of energy is converted into another. The first and foremost constraint on energy transformation is that total energy must be conserved. If we account properly for all energy conversion and energy transfer processes, the total amount of energy present must remain constant. To account properly for all types of energy, we will need to define a number of terms quite carefully. First, we must specify precisely what is being studied. The system is defined as the part of the universe that is being considered. The remainder of the universe is referred to as the surroundings, even though it is not generally necessary to consider everything else in the actual universe. These definitions assure that the system plus the surroundings must equal the universe. The system and the surroundings are separated by a boundary. In some cases, this boundary may be a physical container, and in others, it might be a more abstract separation. These ideas may seem obvious, but the choice of system and surroundings is not always clear. For example, if the system to be studied is “the atmosphere of the earth,” the definition of the boundary is somewhat challenging. The atmosphere simply gets gradually less dense, and we must make some arbitrary decision as to where it ends. In most cases, the main requirement is that we be consistent. The same choice of system and surroundings must be used throughout a particular problem, even if the choice is somewhat arbitrary. Once an appropriate choice of a system has been made, the concept of conservation of energy immediately becomes useful. Because we said that heat and work


are the only possible forms of energy transfer, we can attribute the overall change in energy, E, of a system to these two components. Heat is commonly designated as q and work as w, so we can write DE = q + w

(9.2)

The symbol D (delta) is introduced here as a notation meaning “the change in.” This symbol, which will be used frequently in our study of thermodynamics, is always defined as the difference between the final state and the initial state. DE = Efinal – Einitial

(9.3)

Equation 9.2 is deceptively simple because it includes arbitrary choices for the meaning of the signs for the quantities of heat and work. Again, the key is to choose consistent definitions for those signs. Convention dictates that energy transferred into a system is given a positive sign and energy flowing out of a system carries a negative sign. Thus when heat flows into a system from the surroundings, the value of q is positive, and when work is done on a system, the value of w is positive. Conversely, when heat flows out of a system or work is done by the system on the surroundings, q and w will be negative. Example Problem 9.1 demonstrates the use of this equation.


9 .1

If 515 J of heat is added to a gas that does 218 J of work as a result, what is the change in the energy of the system?

Strategy Energy flow occurs as either heat or work, and Equation 9.2 relates those to the total energy. We are given the magnitudes for q and w, but we must determine the direction for each process to assign the correct signs.

Solution Heat added TO the system means that q > 0, so q = +515 J. Work done BY the system means that w < 0, so w = –218 J. DE = q + w = 515 J + (–218 J) = +297 J Note that in most cases when a value is positive, the “+” sign is not placed in front of the number.

Discussion Though fairly simple numerically, this problem points to the need to consider the signs of q and w carefully.

Check Your Understanding 408 J of work is done on a system that releases 185 J of heat. What is the energy change in the system?

In the example above, we see that DE is not zero. How can this be consistent with the idea that energy must be conserved? By following our sign conventions for q and w, we considered the processes of heat and work from the perspective of the system (the unspecified gas in this case). So the value of DE we calculated is also the change in internal energy for the system. This change in the internal energy of the system is precisely offset by a change in the surroundings: DEsurroundings = –DEsystem. (Usually, we just write DE rather than DEsystem because, by definition, the system is what we are most interested in.) Thus the energy of the universe remains constant. DEuniverse = DEsurroundings + DEsystem = 0 Energy can be transformed from one form to another but cannot be created or destroyed. This is known as the first law of thermodynamics.

9.3


287

Waste Energy

Of course, on a cold winter day, some of the waste heat from the engine is used to warm the passenger cabin. So clever engineering can put “wasted” heat to use.

We have already alluded to the conversion of energy from one form to another a number of times in this chapter. Anyone who sits in front of a campfire to ward off the evening chill can argue the virtues of heat, but in most cases, civilization prefers its energy in the form of work. The combustion of gasoline is not inherently useful, but when the heat released is harnessed in the engine of an automobile, the resulting work gets us where we need to go. All available observations, however, point to the idea that it is impossible to convert heat completely to work. (This leads to the second law of thermodynamics, which we’ll examine in detail in Chapter 10.) The car’s engine gets hot when it runs. The heat that warms the engine does not propel the car toward its destination. So a portion of the energy released by the combustion of gasoline does not contribute to the desired work of moving the car. In terms of the energy economy, this energy can be considered wasted. One common way to obtain work from a system is to heat it: heat flows into the system and the system does work. But in practice, the amount of heat flow will always exceed the amount of useful work achieved. The excess heat may contribute to thermal pollution. (Thermal pollution is the raising or lowering of water temperature in streams, lakes, or oceans above or below normal seasonal ranges from the discharge of hot or cold waste streams into the water.) The efficiency of conversion from heat to work can be expressed as a percentage. Typical efficiencies for some common conversion processes are shown in Table 9.1. The prospect of improving the efficiencies of appliances is often discussed Table

❚

9.1

Typical efficiencies of some common energy conversion devices

288

Device

Energy Conversion

Electric heater

Electrical : thermal

100

Hair drier

Electrical : thermal

100

Electric generator

Mechanical : electrical

95

Electric motor (large)

Electrical : mechanical

90

Battery

Chemical : electrical

90

Steam boiler (power plant)

Chemical : thermal

85

Home gas furnace

Chemical : thermal

85

Home oil furnace

Chemical : thermal

65

Electric motor (small)

Electrical : mechanical

65

Home coal furnace

Chemical : thermal

55

Steam turbine

Thermal : mechanical

45

Gas turbine (aircraft)

Chemical : mechanical

35

Gas turbine (industrial)


30

Automobile engine


25

Fluorescent lamp

Electrical : light

20

Silicon solar cell

Solar : electrical

15

Steam locomotive


10

Incandescent lamp

Electrical : light


Typical Efficiency (%)

5

300 2009 Technology Best available

200

150

100

50

0 Lighting

Central air conditioners

New building shell efficiency

Refrigerators

Natural gas furnaces

Data are from the Energy Information Administration’s Annual Energy Outlook, 2009.

Efficiency (% gain over 2007)

250

Figure 9.5 ❚ The graph shows the predicted efficiency gains by the year 2030 for various technologies.

as an important means of conserving energy in the future. The possible impact of this is shown in Figure 9.5, which compares two possible scenarios for the efficiency of technologies in the year 2030. In one case, no new technologies are developed after 2009, but the appliances in use gradually become more efficient as people replace older models. In the second case, it is assumed that consumers always choose the best available technology, independent of cost and other factors. But even with savings in energy associated with efficient appliances, the generation of electricity will still always result in energy being lost as heat. Energy output in the form of electrical energy is always less than the chemical or nuclear energy input at the generating plant.

9.4

Heat Capacity and Calorimetry

General statements about nationwide or worldwide energy resources and their consequences may seem abstract in many respects, if only because the numbers are so large. Making the connection between this global scale and systems we can observe in a laboratory requires a systematic way to measure energy flow. We can do this by observing heat flow into or out of a system through a set of techniques collectively called calorimetry.

Heat Capacity and Specific Heat Suppose that we wish to raise the temperature of two different systems or objects. In general, the different systems will absorb different amounts of energy based on three main factors: the amount of material, the type of material, and the temperature change. One way to note the importance of the amount of material is to compare the behavior of a glass of water and an ocean. On a hot summer day at the beach, a glass of cold water will quickly become warm, whereas the ocean water temperature does not change noticeably over the same time. The small amount of water in the glass behaves differently from the large amount in the ocean. The type of material is also important, and the summer beach provides some insight here, too. The sand at most beaches is predominantly silicon dioxide, and it heats up much more quickly 9.4 Heat Capacity and Calorimetry

289

than even the shallowest water. Both materials are exposed to similar amounts of energy from the sunlight, but they behave differently. Finally the amount of energy supplied and the temperature change are also related. On a cloudy day at the beach, the sand does not get as hot: less energy is supplied to the sand because the clouds absorb some energy from the Sun. Summarizing these observations, if we want to calculate the heat associated with a given temperature change, we’ll need to account for the amount and identity of the material being heated as well as the extent of the temperature change. This idea can easily be expressed as an equation, and we will use two slightly different versions of this relationship. We have two options for expressing the amount of material: by mass or by moles. Either choice leads to a usable equation. You already may be familiar with the use of mass because that is commonly done in physics or physical science classes. In that case, the identity of the substance is included through a term called the specific heat capacity (c, usually simply called the specific heat), leading to Equation 9.4. q = mcDT

(9.4)

The specific heat is a physical property of a material that measures how much heat is required to raise the temperature of one gram of that material by 1°C. Similarly, the molar heat capacity is a physical property that describes how much heat is required to raise the temperature of one mole of a substance by 1°C. So if we choose to express the amount of material in terms of moles rather than mass, our equation changes only slightly. q = nCpDT

(9.5)

(The subscript p on Cp indicates that this is the heat capacity at constant pressure. Under other conditions, such as constant volume, the value of the heat capacity may differ slightly.) Either equation provides the same information about the heat needed to produce a given change in temperature. As long as we know the molar mass of the substance, it should be simple to convert between the specific heat and the molar heat capacity. Table 9.2 provides a list of specific heats and molar heat capacities for a few materials. A more extensive table can be found in Appendix D. Example Problems 9.2 and 9.3 show applications of these equations. As you work through these examples, notice that the heat capacity will have the same numerical value whether the temperature is expressed in K or in °C. Although this may seem odd at first, it makes sense when you realize that the heat depends only on the temperature difference, not on the temperature itself. Verify for yourself that if the temperature of a piece of metal rises from 20 to 30°C, DT can correctly be written as either 10°C or 10 K.

Table

❚

9.2

Specific heat and molar heat capacities for some common substances

Specific Heat, c ( J g–1 K–1)

Molar Heat Capacity, Cp ( J mol–1 K–1)

Al(s)

0.900

24.3

Cu(s)

0.385

24.5

H2O(s)

2.09

37.7

H2O(ℓ)

4.18

75.3

H2O(g)

2.03

36.4

Substance Note that water has a much higher heat capacity than the metals shown in this table. This is true for most metals.

290



9.2

Heating a 24.0-g aluminum can raises its temperature by 15.0°C. Find the value of q for the can.

Strategy Heat flow depends on the heat capacity of the material being heated, the size of the sample, and the change in temperature. Because the information given is in mass, it will be easiest to use the specific heat rather than the molar heat capacity. The value of the specific heat of aluminum can be found in Table 9.2.

Solution q = mcDT 0.900 J = 24.0 g × ———— × 15.0˚C g °C = 324 J

Analyze Your Answer It may take some practice to gain a feel for the correct magnitude for heat. We do know that q should have energy units, and the units in our calculation are consistent with that. We can also look at the order of magnitude for the quantities involved to make sure our arithmetic makes sense. The mass and the temperature change are both on the order of 101, whereas the specific heat is on the order of 100, or 1. So the product of the three should be on the order of 102, and an answer in hundreds of joules makes sense.

Check Your Understanding A block of iron weighing 207. g absorbs 1.50 kJ of heat. What is the change in the temperature of the iron?


9.3

The molar heat capacity of liquid water is 75.3 J/mol K. If 37.5 g of water is cooled from 42.0 to 7.0°C, what is q for the water?

Strategy The heat flow is proportional to the amount of water and the temperature change. As before, we can do the calculation with the amount of water expressed in either grams or moles. Molar heat capacity is given but the amount of water is provided in grams. So we will use the molar mass to convert the amount from grams to moles. (We could also leave the amount in grams and convert the molar heat capacity into the specific heat.) Care must also be taken in defining the change in temperature. Such changes are always the final state minus the initial state. Because the water is cooling here, DT will be negative.

Solution q = nCpDT 75.3 J 1 mol = 37.5 g × ——— × ———— × – 35.0°C 18.0 g mol °C = –5.49 × 103 J = –5.49 kJ The negative value indicates that the system (in this case, the water) has lost energy to the surroundings. Notice that as long as we correctly express DT as Tfinal – Tinitial, the correct sign for q will result automatically.

9.4 Heat Capacity and Calorimetry

291

Check Your Understanding If 226 kJ of heat increases the temperature of 47.0 kg of copper by 12.5°C, what is the molar heat capacity of copper?

In these examples, we have simply stated that a given amount of heat was added to or removed from our system, without considering the source of that heat. Although this lets us familiarize ourselves with the equations, it isn’t very realistic. In a more common application, we can use the same type of equations to determine heat flow between two objects, as shown in Example Problem 9.4.

E XAM PL E PRO BLEM

9 .4

A glass contains 250.0 g of warm water at 78.0°C. A piece of gold at 2.30°C is placed in the water. The final temperature reached by this system is 76.9°C. What was the mass of gold? The specific heat of water is 4.184 J/g °C, and that of gold is 0.129 J/g °C.

Strategy We will need to assume that heat flows only between the gold and the water, with no heat lost to or gained from the glass or the surroundings. Then the heats must balance; the heat gained by the gold was lost by the water. Each of those can be calculated using Equation 9.3, and the two must be equal and opposite in sign.

Solution qgold = –qwater mgold × cgold × DTgold = – mwater × cwater × DTwater 0.129 J 4.184 J mgold × ———— × 74.6°C = –250.0 g × ———— 3 –1.1°C g °C g °C Rearranging gives us 4.184 J −250.0 g × ———— × −1.1°C g °C

mgold = —————————————— = 120 g 0.129 J ———— × 74.6°C g °C

Analyze Your Answer From the original problem statement, we can see that the temperature of the hot water doesn’t change very much. This small DT for the hot water implies that the gold sample must be fairly small. Next, consider that the heat capacity for water is roughly 40 times larger than that for gold, and the temperature change for gold is almost 80 times greater than for water. These two observations suggest that the gold sample should be close to half the size of the water sample, and our calculated result confirms this.

Check Your Understanding A 125-g sample of cold water and a 283-g sample of hot water are mixed in an insulated thermos bottle and allowed to equilibrate. If the initial temperature of the cold water is 3.0°C, and the initial temperature of the hot water is 91.0°C, what will be the final temperature?

These examples provide a means to calculate heat flow or related information. But we have yet to see how we can actually make laboratory measurements of heat flow.

292


Calorimetry Calorimetry is the term used to describe the measurement of heat flow. Experiments are carried out in devices called calorimeters that are based on the concept established in Example Problem 9.4. The heat evolved or absorbed by the system of interest is determined by measuring the temperature change in its surroundings. Every effort is made to isolate the calorimeter thermally, preventing heat flow between the immediate surroundings and the rest of the universe. Figure 9.6 shows a typical calorimeter. If the instrument is thermally isolated from the rest of the universe, the only heat flow that must be considered is that between the system being studied and the immediate surroundings, whose temperature can be measured. A two-step process is used to make a calorimetric measurement. The first step is calibration in which a known amount of heat is generated in the apparatus. The second step is the actual measurement, in which we determine the amount of heat absorbed or released in the reaction of a known amount of material. The calibration can be done either by burning a known amount of a well-characterized material or by resistive heating, in which a known amount of current is passed through a wire that heats due to its electrical resistance. The heat capacity of the entire calorimeter may be obtained by measuring the change in temperature of the surroundings resulting from a known heat input.

There are many specialized types of calorimetry, but all involve this sort of calibration.

Known amount of heat = calorimeter constant × DT q = Ccalorimeter × DT

or

(9.6)

Note that in contrast to our earlier equations relating q and DT, there is no mass or number of moles term here for the quantity of material. The calorimeter constant is the heat capacity of a particular object (or set of objects) rather than that of a material. It may help to think of it as the heat capacity “per calorimeter” and then realize that we have just one calorimeter. For someone who routinely uses the same calorimeter, this approach is much simpler than the alternative, which would be to keep track of the masses of steel, water, and other materials in the calorimeter. In the case of a bomb calorimeter such as that shown in Figure 9.6, the calorimeter constant is largely attributable to the water that surrounds the bomb but also includes the heat capacities of the thermometer, the stirring system, and the bomb itself. Once the calorimeter constant is known, we are ready to use the calorimeter for our actual measurement. We place known amounts of reactant(s) into the calorimeter,

Ignition wires heat sample.

Figure 9.6 ❚ A bomb calorimeter

Thermometer Stirrer

Surroundings Universe System

Water

Heat

No heat flow –qsys ⫽ ⫹qsurr Insulated outside chamber

Sample dish

Burning sample

Steel bomb

is a fairly complicated piece of equipment, as shown in the diagram on the left. But the general premise of the device is simply to carry out a reaction at constant volume and with no heat flow between the calorimeter and the outside world. The diagram on the right shows the standard choice of system and surroundings for a bomb calorimetry experiment. The system consists of the contents of the bomb itself. The surroundings include the bomb and the water bath surrounding it. We assume that no heat is exchanged with the rest of the universe outside the insulated walls of the apparatus.

9.4 Heat Capacity and Calorimetry

293

initiate the reaction, and then measure the resulting temperature change of the calorimeter. The calorimeter constant allows us to determine the amount of heat released or absorbed in the reaction. Example Problem 9.5 shows how this process works.

E XAM PL E PRO BLEM

We will explore the idea of energy density further in Section 9.7.

9 .5

A calorimeter is to be used to compare the energy content of some fuels. In the calibration of the calorimeter, an electrical resistance heater supplies 100.0 J of heat and a temperature increase of 0.850°C is observed. Then 0.245 g of a particular fuel is burned in this same calorimeter, and the temperature increases by 5.23°C. Calculate the energy density of this fuel, which is the amount of energy liberated per gram of fuel burned.

Strategy The calibration step allows us to determine the calorimeter constant. Once this is known, the amount of heat evolved from the fuel can be determined by using Equation 9.6. Finally, we divide this heat by the mass of fuel that generated it to arrive at the requested energy density.

Solution Step 1: Calibration q = Ccalorimeter × DT so Ccalorimeter = q/DT = 100.0 J/0.850°C Ccalorimeter = 118 J/°C Step 2: Determination of heat evolved by fuel qcalorimeter = Ccalorimeter × DT = 118 J/°C × 5.23°C = 615 J And qfuel = –qcalorimeter = –615 J Step 3: Calculation of the energy density Energy density = –qfuel/m = –(–615 J)/0.245 g = 2510 J/g = 2.51 kJ/g

Discussion This problem illustrates the need to be careful with signs in thermodynamic calculations. Because the burning of fuel releases heat, q for the fuel should be negative. The energy density, though, would be reported as a positive number, resulting in the additional negative sign in the final step.

Check Your Understanding The combustion of naphthalene (C10H8), which releases 5150.1 kJ/mol, is often used to calibrate calorimeters. A 1.05-g sample of naphthalene is burned in a calorimeter, producing a temperature rise of 3.86°C. Burning a 1.83-g sample of coal in the same calorimeter causes a temperature change of 4.90°C. What is the energy density of the coal?

294


9.5

Enthalpy

The concept of heat flow is reasonably intuitive. Cool objects get warmer at the expense of hotter objects. Scientifically, however, this natural process requires more careful examination. The conditions under which heat flow occurs will have an impact on the measurements that are made. For example, when one mole of octane, C8H18, burns to form gaseous carbon dioxide and liquid water under constant volume conditions, 5.45 × 103 kJ of heat is released. The same reaction under constant pressure conditions releases 5.48 × 103 kJ. The percentage difference between the two conditions is small, but consider the engineering consequences in designing safety systems for large octane storage containers. Such vessels may hold in excess of 1000 moles of octane. Each mole releases 30 kJ more energy under combustion conditions of constant pressure versus constant volume. When thousands of moles burn, the difference in total energy release may be important for safety. Why do the two cases release different amounts of heat, and how can we account for these differences properly?

Defining Enthalpy Because of the observation that heats of reaction depend on reaction conditions, it has proven advantageous to define two mathematical functions for energy. Internal energy has already been introduced. By using the definition of internal energy, we can show that under constant volume conditions, the change in internal energy equals the heat flow. Start with DE = q + w In chemical reactions, we usually need to consider only PV-work. When a gas expands, it does an amount of work equal to PDV on its surroundings. But if the expanding gas is our system, we want w to be the work done on the gas, and that will be –PDV. So we can replace w in the equation above with –PDV. DE = q – PDV

(9.7)

This equation now allows us to see how internal energy is related to heat. If the volume is held constant, DV is zero, so the second term is zero. All that remains is DE = qv

(9.8)

where the subscript “v” is added to denote that the equation is correct only under constant volume conditions. Equation 9.8 tells us that if we do a calorimetry experiment at constant volume, we will get a direct determination of DE. But what if we do the experiment at constant pressure instead? It would be useful to have a function that would be equal to the heat flow under constant pressure conditions. This function, known as enthalpy, is defined as H = E + PV

(9.9)

Working from this definition, we can show that the change in enthalpy (DH) will be equal to the heat flow under constant pressure conditions. From the above definition, the change in enthalpy (DH) must be DH = DE + D(PV) We can expand this by substituting for DE using Equation 9.7. DH = (q – PDV ) + D(PV ) If the pressure is held constant, then the D(PV ) term will simply become PDV, giving DH = q – PDV + PDV The second and third terms clearly cancel and leave the desired result. DH = qp

(9.10) 9.5

Enthalpy

295

The enthalpy change therefore equals the heat flow under constant pressure. (This is denoted with a subscript “p,” just as we did in Equation 9.8 for constant volume.) Now we have two ways to define heat flow into a system, under two different sets of conditions. For a process at constant volume, the measurable heat flow is equal to DE, the change in internal energy. For a process at constant pressure, the measurable heat flow is equal to the change in enthalpy, DH. In many ways, enthalpy is the more useful term because constant pressure conditions are more common. A reaction carried out in a beaker in the chemistry laboratory, for instance, occurs under constant pressure conditions (or very nearly so). Thus, when we refer to the heat of a process, we are typically referring to a change in enthalpy, DH. As in previous definitions, DH refers to Hfinal – Hinitial. When heat evolves from a system, the process is said to be exothermic and the value of DH is less than zero. An exothermic process feels hot: if you pick up the beaker in which an exothermic reaction is taking place, heat will flow from the reacting system into your hand. Conversely, when heat is absorbed by the system, the process is said to be endothermic, and the value of DH is greater than zero. Endothermic processes feel cold because they draw heat from their surroundings.

DH of Phase Changes

The term “fusion” here is synonymous with melting.

For any substance, the heat of fusion is typically much smaller than the heat of vaporization.

Heat flow into a substance does not always raise its temperature. If heat flows into an ice cube at 0°C, for example, the ice will melt to form liquid water at 0°C. (If heat continues to flow into the resulting water, its temperature will begin to rise, of course.) How can the temperature remain constant despite the influx of heat? To understand this, we need to remember that intermolecular forces are more extensive in a solid than in a liquid. So as the ice cube melts, energy must be expended to overcome some of the intermolecular attractions. The internal energy of liquid water is higher than that of solid ice, even though both are at the same temperature. For similar reasons, there will be heat flow in any phase change. The names of phase changes among solids, liquids, and gases are summarized in Figure 9.7. Because these phase changes generally take place at constant pressure, the corresponding heat flows should be viewed as changes in enthalpy. Some phase changes are so common that their enthalpy changes have specific names and symbols assigned to them. The heat required to melt a substance is the heat of fusion, DHfus. The enthalpy change for converting a liquid to a gas is known as the heat of vaporization, DHvap. We know that when a liquid is converted to a gas, all the molecules in a sample must overcome whatever intermolecular forces are present. That means that energy must flow into the substance to vaporize it, so the heat of vaporization will always be positive. It follows that the reverse process, condensation, will always release heat. The values of enthalpy changes in opposite directions have the same numeric value and differ only in their signs. Because the strength of the intermolecular forces varies from one substance to another, the magnitude of the enthalpy change for any phase transition will also depend on the substance involved. Values can be found in many standard tables, including one in Appendix D. Data for water are provided in Table 9.3 as an example. Table

❚

9.3

Standard molar enthalpies and temperatures for phase changes of water

Phase Change Transition temperature DH (J/mol)

296


Fusion

Freezing

Vaporization

Condensation

0ºC

0ºC

100ºC

100ºC

DHfus = 6009.5

DHfreeze = –6009.5

DHvap = 4.07 × 104

DHcond = –4.07 × 104

Figure 9.7 ❚ Microscopic level views of solid, liquid, and gas phases are shown, along with the terms commonly used for the possible transitions among these phases.

ion sat en ion nd rat Co apo Ev

De po siti Su on bli ma tio n

Gas

Melting Solid

Freezing

Liquid

These data would allow us to calculate the heat required for a phase transition involving any particular sample of water. Just as we saw when we considered temperature changes, the heat will depend on the amount of substance undergoing the transition: it will take more heat to melt a large block of ice than a small cube. The values shown in Table 9.3 are given in units of J/mol, so we might want to specify the amount of water in terms of moles, leading to the following equation: DH = n × DHphase change

(9.11)

Here n is the number of moles, as usual. Notice that this relationship does not include DT, in contrast to Equations 9.4 and 9.5. This should make sense if we keep in mind that phase changes occur at a constant temperature, so there is no DT in the transition.


9.6

Calculate the enthalpy change when 240. g of ice melts.

Strategy The DHfus value in Table 9.3 is in J/mol, so the amount of ice must be converted into moles. Multiplying the number of moles by DHfus will provide the desired quantity.

Solution 1 mol 240 g H2O × ——— = 13.3 mol H2O 18.0 g DH = n × DHfus = 13.3 mol × 6009.5 J/mol = 8.01 × 104 J

Analyze Your Answer The enthalpy of fusion is constant at 6 kJ/mol, so the enthalpy change depends on the size of the sample. Because the molar mass of water is 9.5

Enthalpy

297

18 g/mol, a sample that is 240 grams is a bit more than 10 moles. That means we should expect an answer that is a bit larger than 60 kJ, which is consistent with our result.

Check Your Understanding Calculate the enthalpy change when 14.5 g of water vapor condenses to liquid water.

Just as for a temperature change, there is no fundamental reason why we could not express the quantity of material in mass rather than moles. But it is essential that we be consistent in the choice of units: if the quantity is in mass, then we will need an enthalpy change value that is expressed in mass as well. Because you know the molar mass of water, it should be easy to convert the values in Table 9.3 from J/mol to J/g. If you repeat the calculations from the last example that way, you should get the same result. In most engineering fields, it is more common to use mass rather than moles to specify the amount of a substance. But the underlying concept is the same no matter which units we use. At this point we can combine what we’ve learned about heat flow for both phase changes and temperature changes to obtain the enthalpy change for converting water from ice to liquid and then to steam. A plot of temperature versus heat flow for such a process is shown in Figure 9.8. The heat transfers in sections (1), (3), and (5) in this figure are for changes of temperature, and sections (2) and (4) represent phase changes.

Vaporization and Electricity Production

It is interesting to wonder how the design of power plants would change if something with a much smaller heat of vaporization replaced water.

The large amount of energy needed for converting water from a liquid to a gas, as by the series of processes shown in Figure 9.8, is exploited in converting chemical energy into electricity. The basic features of a fossil fuel–powered electricity plant are shown schematically in Figure 9.9. When the fuel—typically coal or natural gas—burns, chemical energy is released as heat. The goal of the power plant is to convert as much of this energy as possible into electricity. The critical step in this process is to trap the heat given off in the combustion reaction. Water is the material of choice for this process because it has a large heat of vaporization. You’ll recall from Section 8.4 that intermolecular forces are unusually strong in water because of extensive hydrogen bonding. These hydrogen bonds between water molecules in the liquid lead to the large value of DHvap. If a material with a small heat of vaporization were substituted for water, much more of ⫹200 Heat liberated

5

Temperature °C

⫹150 4 Boiling

⫹100

STEAM

3

⫹50

LIQUID WATER (0°–100°C)

Melting 2

0

1 ⫺50

(100°–200°C)

0

Heat absorbed

ICE (⫺50°– 0°C)

200

400

600

800 1000 Heat (kJ)

1200

1400

1600

Figure 9.8 ❚ The figure shows how the temperature of a 500-g sample of H2O varies as it absorbs heat. Here we begin with ice at –50°C. In the segment of the graph marked (1), the temperature of the ice rises until it reaches the melting point at 0°C. In region (2), ice melts at a constant temperature of 0°C until it is converted into liquid. In region (3), only liquid water is present, and its temperature rises until it reaches the boiling point at 100°C. In region (4), the water boils at a constant temperature of 100°C until it is converted into steam. Finally, in region (5), the resulting steam continues to absorb heat, and its temperature rises.

298


Figure 9.9 ❚ The important To power grid

Steam

Steam turbine (generates electricity)

elements of a standard electric power plant are shown in this schematic diagram. The process exploits the large heat of vaporization of water.

Condenser (steam from turbine is condensed by river water)

Steam generator

Pump P P Pump Burning fossil fuels

80° F

100° F

River

it would be required to absorb the same amount of heat. The large heat of vaporization coupled with the relative abundance of water has led to its widespread use in the energy sector of the economy.

Heat of Reaction So far we have considered enthalpy changes for simple physical processes such as temperature changes and phase transitions. But the importance of chemistry to the energy economy arises from the fact that there are enthalpy changes in chemical reactions as well. This enthalpy change is commonly referred to as the heat of reaction. Because many reactions are carried out under constant pressure conditions, this term is sensible, even if slightly imprecise.

Bonds and Energy Chemical reactions involve energy changes because chemical bonds are broken and formed when reactants are transformed into products. Consider a fairly simple reaction—the combustion of methane. CH4(g) + 2 O2(g) : CO2(g) + 2 H2O(ℓ) We are familiar with all of these molecules, so we can easily write Lewis structures for them. H H

C

H⫹ 2O

O

O

C

O⫹2

H

O

H

H On the reactant side of the equation, we have four C!H single bonds and two O"O double bonds. On the product side, we have two C"O double bonds and four O!H single bonds. Over the course of the reaction, all bonds in the reactants must be broken, and that will require an input of energy. On the other hand, all bonds in the products must be formed, and that will release energy. (Recall that bond formation is always exothermic and bond breaking is always endothermic.) If the energy released in forming new bonds is greater than the energy required to break the original bonds, then the overall reaction should be exothermic. Conversely, if the energy needed for

9.5

Enthalpy

299

bond breaking is greater than that released in bond making, then the reaction will be endothermic. Because methane is the principal component of natural gas, we know that burning it must release energy. So we should anticipate that this particular reaction must be exothermic, and calorimetric measurements bear this out; DH for the combustion of one mole of methane is –890.4 kJ. We might attempt to use estimates for the individual bond energies to calculate an approximate value of DH for a reaction. This is sometimes done to obtain an estimate for reactions involving compounds where no thermochemical data are available. The accuracy of this approach is not very good, however, because tabulated values for bond energies represent averages. As we will see in the following section, a simple and accurate alternative exists in most cases. The overall energetics of a chemical reaction can be summarized in a thermochemical equation. For the combustion of methane, the thermochemical equation is CH4(g) + 2 O2(g) : CO2(g) + 2 H2O(ℓ)

DH = –890.4 kJ

This equation tells us two important things. First, we can immediately see whether the reaction is exothermic or endothermic because DH carries a sign. We know that the combustion of methane must be exothermic, and the negative sign on DH confirms this. Second, the thermochemical equation includes the numerical value of DH, so we can tell exactly how much heat will be released. It is important to realize that the heat of reaction shown is for the equation exactly as written; if one mole of methane reacts with two moles of oxygen, 890.4 kJ will be released. But if more fuel is burned, more heat will be released. So if the stoichiometric coefficients are multiplied by some factor, the heat of reaction must also be multiplied by that factor. Thus another thermochemical equation for the combustion of methane is 2 CH4(g) + 4 O2(g) : 2 CO2(g) + 4 H2O(ℓ)

DH = –1780.8 kJ

Heats of Reaction for Some Specific Reactions

The “degree sign” appearing on this symbol for the heat of formation indicates that the value is for a process occurring under standard conditions of 25ºC and 1 atm.

Some classes of chemical reactions are sufficiently common or useful that they have been assigned their own label for heats of reaction. The example reaction for methane we have already used falls into one such category, combustion. Because combustion is a common part of the energy economy, enthalpy changes for combustion reactions are used to compare various fuels. Sometimes these heats of combustion are designated as DHcomb. Similarly, the neutralization reactions between acids and bases have heats of reaction that are called heats of neutralization, DHneut. Another class of reactions, formation reactions, has specially designated heats of reaction known as heats of formation, DHf°, even though many of them are never carried out in practice. Their importance is due to the computational utility derived from the following definition. A formation reaction is the chemical reaction by which one mole of a compound is formed from its elements in their standard states. The standard state is the most stable form of the element at room temperature (25°C) and pressure (1 atm). The formation reaction for carbon monoxide is 1 O (g) : CO(g) C(s) + — 2 2

DH ° = DHf °[CO(g)]

Note that the definition of the formation reaction immediately implies that the heat of formation for any element in its standard state must always be zero. To see this, imagine writing the formation reaction for an element in its standard state, such as O2(g). We would need one mole of O2(g) as the product and oxygen in its standard state as the reactant. But because the standard state of oxygen is O2(g), our formation reaction is really no reaction at all. O2(g) : O2(g)

DH ° = DHf °[O2(g)] = 0

Because both sides of this equation are the same, there can be no change in enthalpy. So DH ° must be zero. This will be true for any element in its standard state. 300


Because we require just one mole of CO on the right-hand side, we must use a fractional coefficient on the O2 to balance the equation, even though this may look odd. Two very common errors when writing formation reactions are including elements that are not in their standard states or having more than one mole of product compound. For carbon monoxide, these mistakes might produce either of the following incorrect “formation” reactions. C(s) + O(g) : CO(g)

DH ° ≠ DHf °[CO(g)]

2 C(s) + O2(g) : 2 CO(g)

DH ° ≠ DHf °[CO(g)]

Although both are valid chemical equations, they do not describe the formation reaction for CO. The first reaction is not correct because oxygen is not shown in its standard state, as a diatomic molecule. The second reaction is not correct because two moles of carbon monoxide are formed. Formation reactions are quite useful for determining heats of reaction, as we will see in the following section.

9.6

Hess’s Law and Heats of Reaction

The enthalpy changes in chemical reactions play an important role in a wide range of situations. Comparisons of possible new fuels or explosives are obvious examples. In many cases, it might be desirable to determine the enthalpy change for a reaction without necessarily carrying out the reaction itself. The synthesis of a new explosive might be very difficult and dangerous, for example, so it would help to know in advance if its potential usefulness might justify the effort. In other cases it may simply be very difficult to perform a direct calorimetric determination of a heat of reaction. For a variety of reasons, we often need to obtain heats of reaction indirectly.

Hess’s Law To obtain information about heats of reaction indirectly, we use what is known as Hess’s law: the enthalpy change for any process is independent of the particular way the process is carried out. The underlying concept upon which this idea is built is that enthalpy is a state function. A state function is a variable whose value depends only on the state of the system and not on its history. When you drive your car, your location is a state function, but the distance that you have traveled to get there is not. For a chemical reaction, the concept of state functions can be very important. We rarely if ever know the microscopic details of how reactant molecules are actually converted into product molecules. But it is relatively easy to determine what the reactants and the products are. Because enthalpy is a state function, if we can find a way to determine the change in enthalpy for any particular path that leads from the reactants to the desired products, we will know that the result will apply for whatever actual path the reaction may take. The block diagram in Figure 9.10 illustrates this concept. ΔHdesired

Initial state

ΔHAf

ΔHAi ΔHBi

Final state

Intermediate state A

ΔHBf

Pressure, volume, and temperature are all state functions.

Figure 9.10 ❚ Conceptual diagram representing Hess’s law. Because enthalpy is a state function, we can choose any convenient path from the initial state to the final state and use that path to calculate the enthalpy change.

Intermediate state B

9.6 Hess’s Law and Heats of Reaction

301

Figure 9.11 ❚ Enthalpy diagram CH4 ⫹ 2 O2

Enthalpy

for the combustion of methane. Here we imagine that one mole of CH4 is first converted to CO (step 1) and that the CO then reacts further to form CO2 (step 2). If we know the values of DH1 and DH2, we can use them to calculate DHcomb.

ΔH1 ΔHcomb

CO ⫹ 2 H2O ⫹ 1/2O2 ΔH2 CO2 ⫹ 2 H2O

For the situation shown in Figure 9.10, the desired enthalpy change could be obtained via either of two alternative pathways. DHdesired = DHAi + DHAf DHdesired = DHBi + DHBf We can develop this idea further by specifying a system and including an enthalpy axis in the picture. Figure 9.11 shows such an enthalpy diagram for the combustion of methane. As we saw in Section 4.1, this combustion may either be complete, producing carbon dioxide directly, or incomplete, producing carbon monoxide. (The production of carbon monoxide from burning of methane presents a danger for natural gas-burning furnaces that are not vented properly.) Hess’s law has important practical implications. Many times we can break a chemical reaction down into a series of steps whose net result is the same as that of the original reaction. According to Hess’s law, in such a case, we can evaluate the enthalpy of the original reaction by using the sequential ones.

E XAM PL E PRO BLEM

9 .7

Sulfur trioxide reacts with water to form sulfuric acid, a major contributor to acid rain. One origin of SO3 is the combustion of sulfur, which is present in small quantities in coal, according to the following equation. 3 S(s) + — O2(g) : SO3(g) 2 Given the thermochemical information below, determine the heat of reaction for this reaction. S(s) + O2(g) : SO2(g)

DH° = –296.8 kJ

2 SO2(g) + O2(g) : 2 SO3(g)

DH° = –197.0 kJ

Strategy We need to use the reactions whose enthalpy changes are known to construct a pathway that results in the desired reaction. In this case, we can see that SO2 is formed in the first reaction we are given and then is consumed in the second. It is critical to pay attention to the precise stoichiometry of the given and desired reactions. The reaction of interest produces just one mole of SO3, whereas the second reaction we are given produces two moles. So we will need to correct for that.

Solution We start with the first of the two given reactions. S(s) + O2(g) : SO2(g)

302


DH° = –296.8 kJ

Next we multiply the second given reaction by one-half. This will account for the fact that the desired reaction produces only one mole of SO3. 1 2

— × [2 SO2(g) + O2(g) : 2 SO3(g)

DH° = −197.0 kJ]

This gives an equation in which one mole of SO2 is consumed, along with its enthalpy change. 1 SO2(g) + — O2(g) : SO3(g) 2

Note that the —12 multiplies the stoichiometric coefficients and the enthalpy change.

DH° = −98.5 kJ

Adding this to the first reaction above gives the desired amount of SO3 and the heat of reaction: S(s) + O2(g) : SO2(g)

DH° = −296.8 kJ

1 O (g) : SO (g) SO2(g) + — 2 3 2

DH° = −98.5 kJ

3 S(s) + — O2(g) : SO3(g) 2

DH° = −395.3 kJ

Analyze Your Answer We may not have any intuition regarding the particular chemical reactions involved, so we’ll have to look at the structure of the problem to see if our answer makes sense. Both of the reactions we added are exothermic, so it makes sense that the combination would be more exothermic than either of the individual reactions.

Check Your Understanding Use the following thermochemical equations as needed to find the heat of formation of diamond: C(diamond) + O2(g) : CO2(g)

DH° = –395.4 kJ

2 CO2(g) : 2 CO(g) + O2 (g)

DH° = 566.0 kJ

C(graphite) + O2(g) : CO2(g)

DH° = –393.5 kJ

2 CO(g) : C(graphite) + CO2(g)

DH° = –172.5 kJ

Formation Reactions and Hess’s Law The type of calculation shown in the previous example has occasional application in chemistry. But Hess’s law is also useful in another more common approach. Heats of formation for many substances are widely tabulated (see Appendix E). Hess’s law allows us to use these tabulated values to calculate the enthalpy change for virtually any chemical reaction. The block diagram in Figure 9.12 shows how this useful scheme arises. ΔHdesired

Initial state Step 1

Heats of formation and other thermodynamic data can be found in reference books such as the Handbook of Chemistry and Physics, or in online sources such as the NIST WebBook.

Final state Step 2

ΔH1 Elements in standard states

ΔH2

Figure 9.12 ❚ This conceptual diagram shows how to use tabulated enthalpies of formation to calculate the enthalpy change for a chemical reaction. We imagine that first the reactants are converted to elements in their standard states, and then those elements recombine to form the products. Because enthalpy is a state function, we do not need to know anything about the actual pathway that the reaction follows.

9.6 Hess’s Law and Heats of Reaction

303

Step 1 is the decomposition of reactants into elements in their standard states. But this is just the opposite of the formation reaction of the reactants, so the enthalpy change of the process is –DHf°(reactants). Similarly, Step 2, the formation of the products from elements in their standard states, has an enthalpy change of DH f°(products). Remember, however, that the formation reaction is defined for the generation of one mole of the compound. Consequently, to use tabulated heats of formation we must multiply by the stoichiometric coefficients from the balanced equation to account for the number of moles of reactants consumed or products generated. Taking these factors into account leads to one of the more useful equations in thermochemistry. DH ° = ∑niDHf ° (product)i − ∑nj DHf ° (reactant)j i

j

(9.12)

In this equation we have designated the stoichiometric coefficients by the Greek letter ν. The first summation is over all of the reactants, and the second is over all of the products. Two example problems will show how heats of formation are used in understanding the thermochemistry of reactions useful in producing energy.

E XAM PL E PRO BLEM

9 .8

Use tabulated data to find the heat of combustion of one mole of propane, C3H8, to form gaseous carbon dioxide and liquid water.

Strategy We will need values for the heats of formation of the reactants and products to determine the desired heat of combustion. First, we must write a balanced chemical equation for the process. Then we can use Equation 9.12 to calculate the heat of the reaction (in this case the heat of combustion) by looking up heats of formation in the table in Appendix E. The stoichiometric coefficients needed will be obtained from the balanced equation. Remember that the heat of formation of an element in its standard state—like the O 2 in this equation—will always be zero.

Solution C3H8(g) + 5 O2(g) : 3 CO2(g) + 4 H2O(ℓ) DH° = 3 mol DHf°(CO2) + 4 mol DHf°(H2O) – 1 mol DHf°(C3H8) – 5 mol (0) = 3 mol (–393.5 kJ/mol) + 4 mol (–285.8 kJ/mol) – 1 mol (–103.8 kJ/mol) = –2219.9 kJ

Discussion We use units of kJ/mol for the heat of formation of a substance. But in writing the enthalpy change of a chemical reaction, we will use kJ as our preferred unit, not kJ/mol. The reaction in this example illustrates why we do this. The value we calculated, DH° = –2219.9 kJ, is for a reaction in which one mole of propane reacts with five moles of oxygen to form three moles of carbon dioxide and four moles of water. So if we were to say “–2219.9 kJ/mol,” we would need to specify carefully which substance that “mol” refers to. We choose to write the DH value in kJ, with the understanding that it refers to the reaction as written. This is also dimensionally consistent with Equation 9.12, provided that we treat the stoichiometric coefficients as carrying units of moles. You may see other texts that refer to values as “per mole of reaction.”

Check Your Understanding Use heat of formation data from Appendix E to calculate DH ° for the following reaction: ClO2(g) + O(g) : ClO(g) + O2(g)

304



9.9

Ethanol, C2H5OH, is used to introduce oxygen into some blends of gasoline. It has a heat of combustion of 1366.8 kJ/mol. What is the heat of formation of ethanol?

Strategy We know the relationship between the heat of combustion of a reaction and the heats of formation of the substances involved. We will need a balanced chemical equation for the combustion, and we must recognize that the combustion of ethanol yields the same products as hydrocarbon combustion. We can write the balanced equation and then use Equation 9.12 to determine the desired quantity. In this case, we know the heat of reaction and will be solving for one of the heats of formation.

Solution C2H5OH(ℓ) + 3 O2(g) : 2 CO2(g) + 3 H2O(ℓ)

DH° = –1366.8 kJ

Using Equation 9.12: DH ° = 2 mol DHf°[CO2(g)] + 3 mol DHf °[H2O(ℓ)] – 1 mol DHf°[C2H5OH(ℓ)] – 3 mol DHf°[O2(g)] –1366.8 kJ = 2 mol (–393.5 kJ/mol) + 3 mol (–285.8 kJ/mol) – DHf°[C2H5OH(ℓ)] – 3 mol (0 kJ/mol) Rearranging and solving gives DHf°[C2H5OH(ℓ)] = –277.6 kJ/mol

Analyze Your Answer By now, we have seen enough heats of formation to know that values in the hundreds of kJ/mol seem fairly typical. Notice that it is very important to handle signs carefully when solving this type of problem: the fact that we are adding and subtracting quantities that can be positive or negative allows many opportunities for errors.

Check Your Understanding Incomplete combustion of hydrocarbons leads to the generation of carbon monoxide rather than carbon dioxide. As a result, improperly vented furnaces can poison people who live in an affected building because of the toxicity of CO. Calculate the heat of reaction for the incomplete combustion of methane, CH4(g), to yield liquid water and CO(g).

9.7


The ability to predict the energetic consequences of a chemical reaction is an important skill in chemistry that has many practical applications. Writing a thermochemical equation allows for a treatment of energy that is similar in many ways to the stoichiometry problems we learned to solve in Chapter 4. For an exothermic reaction, we can treat energy as a product, and in an endothermic reaction, energy can be thought of as a reactant. We must keep in mind that the stated value of DH corresponds to the reaction taking place exactly as written, with the indicated numbers of moles of each substance reacting. The emphasis on the importance of the mole as the heart of the stoichiometry problem remains when we solve problems involving the energy of reactions. So if we wanted to calculate the amount of energy released by burning a given

9.7


305

Given mass

Convert using molar mass

Moles of substance

Given volume & molarity

Use balanced thermochemical equation

Energy released or absorbed

Convert using M ⫻ V = moles

Figure 9.13 ❚ This flow chart shows the sequence of steps needed to calculate the amount of energy released or absorbed when a chemical reaction is carried out using a given amount of material.

Remember that the thermochemical equation tells us the heat for the reaction of the specific molar amounts written.

mass or volume of methane, we would start by converting that quantity into moles. Then we could use the balanced thermochemical equation to relate the amount of energy to the number of moles of methane actually burned. In solving stoichiometry problems, we have used the balanced chemical equation to convert from the number of moles of one compound to the number of moles of another. Now we can use the thermochemical equation to convert between the number of moles of a reactant or product and the amount of energy released or absorbed. Figure 9.13 shows this approach schematically. The conversion factors needed to obtain the number of moles are the same as earlier: molar mass, density, gas pressure or volume, etc. Let’s consider, as an example, the reaction between nitrogen gas and oxygen to form nitric oxide. This occurs regularly as a side reaction when hydrocarbons are burned as fuel. The exothermic combustion reaction that powers an automobile uses air rather than pure oxygen, so there is always a large amount of nitrogen present. At the high temperatures produced in a running engine, some of the nitrogen reacts with oxygen to form NO. N2(g) + O2(g) : 2 NO(g)

DH° = 180.5 kJ

The nitric oxide gas formed in this way is an important species in several pollution pathways and is itself an irritant even when present at fairly low levels. For example, NO(g) reacts further with oxygen to form nitrogen dioxide, whose brown color is largely responsible for the dark haze typical of urban smog. Example Problem 9.10 shows how this type of equation can be used to determine heats of reaction of specific quantities of substances.

E XAM PL E PRO BLEM

9 .10

An engine generates 15.7 g of nitric oxide gas during a laboratory test. How much heat was absorbed in producing this NO?

Strategy The thermochemical equation for this reaction is shown above. This equation provides a link to convert the amount of NO formed to the energy absorbed. As in any other stoichiometry problem, we will work with moles of substance, which we

306


can obtain from the given mass. Note that the stated DH ° value is for the production of two moles of NO because two moles appear in the equation.

Solution 1 mol NO 15.7 g NO × ————— = 0.523 mol NO 30.0 g NO 180.5 kJ 0.523 mol NO × ————— = 47.2 kJ 2 mol NO

Analyze Your Answer The thermochemical equation gives us DH for the reaction as written. That means that 180 kJ are absorbed for every two moles of NO formed. The molar mass of NO is very close to 30 g/mol, so the amount formed in the problem is a little more than half a mole. That’s about a quarter of the amount formed in the thermochemical equation. Our answer is roughly a quarter of the DH value from the thermochemical equation as well, so it seems likely to be correct.

Check Your Understanding If 124 kJ of heat is absorbed in a reaction that forms nitric oxide from nitrogen and oxygen, what mass of NO must have been produced? What mass of N2 was consumed? This type of manipulation can provide insight into the relative merits of various fuels as well.

Energy Density and Fuels When considering the economic merits of a particular fuel, several factors should be considered. Among the typical characteristics of a useful fuel are the availability of technology for extracting it, the amount of pollution released by its combustion, and its relative safety. (Because fuels are burned, there is always some danger of unintended or uncontrolled combustion.) From an economic viewpoint the ease of transporting the fuel is a key factor in this consideration. Fuels that are expensive or dangerous to get to the consumer are less attractive than those that can be delivered at more modest cost. Transportation costs for commodities are dictated largely by the mass that is transported. Because of this important cost, one key feature of a fuel is its energy density, the amount of energy that can be released per gram of fuel burned. The energy densities of some fossil fuels are shown in Table 9.4, along with those of some alternative fuels. It’s easy to see why petroleum is such a prominent source of energy. Consider the positive factors for petroleum use: (1) It is a liquid, which makes it easy to transport and deliver to the customer. (2) It is relatively safe. Explosions can occur, but they are preventable using designs that have a modest cost. (3) The products of its combustion are gases. Where a liquid fuel is easy to transport, gaseous combustion products are easy to dispose of. (4) It has a high energy density. In addition to the issues of transportation cost already discussed, this means that automobile engines can be designed to produce relatively large amounts of energy without having the fuel increase the overall weight of the vehicle by very much. Imagine the complications if a tank of gas for your car weighed as much as the car itself. You would use a significant portion of your gas just to carry its own weight, and vehicle safety would be significantly compromised if the mass of the car were increased by a large amount just because the gas tank was full.

9.7

We introduced the idea of energy density in Example Problem 9.5.


307

Table

❚

9.4

Energy densities for a few possible fuels

Fuel

Energy Density (MJ/kg)

Hydrogen

142.0

Methane

55.5

Octane

47.9

Propane

50.3

Aviation gasoline

43.1

Coal, anthracite

31.4

Diesel fuel

45.3

Oil, crude (petroleum)

41.9

Oil, heating

42.5

Gasoline, automotive

45.8

Kerosene

46.3

Wood, oven dry

20.0

INSIGHT INTO

9.8

Batteries

Throughout this chapter, we have concentrated on reactions that absorb or release energy in the form of heat. Although this is the most common manifestation of energy in chemical reactions, there are also many reactions that interconvert between chemical energy and other forms of energy. Some reactions, such as those that occur in glow sticks or in fireflies, release energy as light rather than heat. You exploit reactions that convert chemical energy into electrical energy every time you turn on a battery-powered device. At the heart of any battery is a chemical reaction that releases energy. How can chemical energy be converted directly into electrical energy? And why do some batteries simply die, whereas others can be recharged more or less indefinitely? We’ll take a closer look at the chemistry of batteries—called electrochemistry—in Chapter 13, but now, we can use ideas from thermochemistry to get a glimpse of the answers. Electrical energy arises from moving a charge. In practice, to power any conventional electrical device, we need to move electrons through a wire. So if a chemical reaction is going to produce useful electrical energy, it will need to function as a source of moving electrons. Because electric charge must always be conserved, these electrons must move through a closed cycle. Electrons “released” at one terminal of the battery must be “recaptured” at the opposite terminal to complete this circuit. What sort of chemical reactions would make good candidates for such a process? Oxidation–reduction (or redox) reactions involve the transfer of electrons from one species to another. A battery, then, is really a cleverly engineered application of a redox reaction. Successful battery design requires that electrons flow through an external circuit as they are transferred from one reacting species to another.

308


In a standard alkaline battery, the important chemical reaction is Zn(s) + MnO2(s) + H2O(ℓ) : ZnO(s) + Mn(OH)2(s) How does this reaction involve the transfer of electrons? Look at the zinc and manganese on each side of the equation. If we treat ZnO as if it were an ionic compound, the Zn would need a 2+ charge to balance the 2– charge on the oxygen. So when zinc metal reacts to form ZnO, each atom of zinc that reacts loses two electrons. The same idea applied to the manganese suggests that it goes from a 4+ state in MnO2 to a 2+ state in Mn(OH)2, so each atom gains two electrons. Thus the net transfer of electrons is from zinc to manganese. The design of the battery (as shown in Figure 9.14) physically separates the zinc and manganese dioxide from one another to ensure that the electrons transferred flow through the external circuit as the reaction proceeds. The redox reaction in a battery has its own characteristic thermochemistry, just like the combustion reactions we have already looked at. The energy change of the chemical reaction fixes the voltage of an electrochemical cell. Some batteries, like the one that starts your car, are actually combinations of multiple cells, arranged so that the effective voltage of the battery is increased. Batteries are sometimes classified as either “primary” or “secondary.” A primary battery such as the typical alkaline cell becomes useless once the underlying chemical reaction has run its course. The lifetime of the battery is determined by the amounts of reactants present, so a relatively large D-cell would last longer than a small AA-cell in the same application. The battery dies when the reactants have been converted into products, bringing the reaction to a halt. In practice, the voltage output of a battery usually begins to decrease near the end of its lifetime, and the cell will generally fail before the reactants are completely consumed. A secondary battery is one that can be recharged, allowing for a much longer life cycle. To make a rechargeable battery, we must be able to reverse the redox reaction, converting the products back into reactants. Because we know that the cell reaction must be exothermic to supply energy, we can also see that the reverse reaction must be endothermic. So some external energy source will be needed to “push” the

Positive cover— plated steel

Figure 9.14 ❚ Construction of a

Can—steel

typical alkaline battery is illustrated.

Electrolyte—postassium hydroxide/water Metalized plastic film label Cathode—manganese dioxide, carbon Separator—nonwoven fabric

Metal spur

Anode—powdered zinc

Current collector— brass pin

Inner cell cover—steel Seal—nylon

Metal washer

Negative cover— plated steel

9.8

Batteries

309

Many of these parameters can be measured quantitatively, so an engineer can make a careful comparison before specifying a choice of battery type.

Table

❚

9.5

Comparison of some characteristics of common primary and rechargeable battery types

Primary Battery Comparison Attribute

Zinc-Air

Alkaline

Lithium

Energy density

High

Medium

High

Energy storage

High

Medium

Medium

Cost

Low

Low

High

Safety

High

High

Medium

Environment

High

High

Medium

Rechargeable Battery Comparison Attribute

Ongoing battery research will continue to have major technological and economic impacts.

Nickel-Cadmium

Nickel Metal Hydride

Lithium Ion

Energy density

Low

Medium

High

Energy storage

Low

Medium

Medium

Cycle life

High

High

High

Cost

Low

Medium

High

Safety

High

High

Medium

Environment

Low

Medium

Medium

reaction back toward the reactants. This is the role of a battery charger, which typically uses electrical energy from some other source to drive the redox reaction in the energetically “uphill” direction. The question of whether or not a given reaction can be made reversible in this way determines whether a particular type of battery can be recharged easily. Many new types of batteries have been introduced in recent years, as demand for portable electronic devices such as cell phones, music players, and laptop computers has soared. What factors influence the choice of battery types for a particular application? One of the most important considerations is energy density, just as we discussed earlier for fuels. Table 9.5 summarizes typical characteristics of several types of batteries. Battery manufacturers usually report energy densities in units of W h/kg. Because a watt is 1 J/s and an hour is 3600 s, 1 W h = 3600 J. Cost also plays a key role in the selection of a battery type. Although some people might be willing to pay hundreds of dollars for a good battery for an expensive laptop computer, you are not likely to pay that much for a battery for your iPod® shuffle. Battery lifetime is also a factor. Digital cameras tend to require more power than traditional film cameras, for example. So consumers are much more willing to pay more for rechargeable batteries for digital cameras because they anticipate cost savings over time.

FOCU S O N

PRO BLEM S O LVI NG

Question During the nuclear reactor accident at Three Mile Island in 1979, an unknown mass of fuel pellets melted, allowing some of the fuel material to fall into water at the bottom of the reactor. Assume that the melting fuel pellets were pure UO2 and had resolidified and cooled to 900°C before reaching the water. Further assume that the water was initially at 8°C and that sensors indicated the final temperature of the water was 85°C. 310


(a) What information would you have to look up to determine the mass of fuel pellets that fell into the water? (b) What would you have to know to use this information to determine the percentage of the fuel that melted in the accident? How would you calculate the percentage?

Strategy (a) The first question that must be addressed is, “What sources of heat raised the temperature of the water?” Were the fuel pellets the only thing that melted? Because we have no other information, we will assume that the only important heat flow was between the pellets and the water. If that were not true, then we would need to account for other materials in our handling of this problem. If the assumption turns out to be incorrect, we will overestimate the amount of fuel that melted in the accident. Once we have made the assumption that the melted fuel is the only source of heat for the water, we would set up a heat balance between the water and the fuel pellets. Thus qpellets = –qwater and (mcDT )pellets = –(mcDT )water We can see there are six possible variables in this equation. Which do we know and which can be looked up? Both DT values are known. The initial temperatures of both the water and the fuel pellets were given in the problem and both share the final temperature of 85°C. The specific heat (c) is not given for either the water or the fuel pellets. This is information that must be looked up, but it should be available from standard reference sources. That leaves the two masses, and clearly the goal of the exercise is to find the mass of the melted fuel pellets. But we don’t know the mass of the water involved, either. So either we must find some way to look up this value or we must make a reasonable estimate. (b) This part of the question assumes that we can find the mass from part (a). If that’s true, then we would need to know the total mass of fuel pellets initially present. This number must be looked up or estimated based on some reliable information.

Answer (a) We would have to look up the heat capacity of water and fuel pellets and the mass of water in the reactor (initially at 8°C) to be able to answer this question. (b) The percentage melted can be calculated using the equation, mmelted Percent melted = ——— × 100% m total

Here we are assuming that we have obtained values for both the mass melted and the total mass of fuel pellets. We should also keep in mind that our result will only be as good as the assumption that the melted fuel was the sole source of heat for the water.

S U M M ARY Energy plays a central role in science and also in economic development and societal advancement. Many energy-related technologies rely on the fact that chemical reactions can release the energy stored in chemical bonds. The study of chemical thermodynamics examines these energy changes in chemical reactions.

In developing a scientific understanding of energy, we must be especially careful to define our terms precisely. For words such as work and heat, the scientific meanings are much more specific than everyday meanings. In scientific usage, work is the movement of a mass against an opposing force, and heat is the transfer Summary

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of energy between objects at different temperatures. As for any other calculations, we also must take care to use consistent units when working with energy values. One familiar concept is that energy is conserved, which is a concise statement of the first law of thermodynamics. To use energy, however, it must often be converted from one form to another. This conversion invariably involves some waste of energy, though, and has a significant impact on the overall energy economy. Experimental measurements of the energy changes in chemical reactions are made by calorimetry, the study of heat flow. Heat flow is important in engineering design, and the ability to relate heat flow to temperature change, for example, is a key skill in both thermochemistry and engineering.

The thermodynamics of chemical reactions is often described in terms of enthalpy. Enthalpy is a state function, meaning that its value depends only on the current state of a system. The enthalpy change is equal to the heat flow for a process at constant pressure. Enthalpy changes for some specific processes, including phase changes, are tabulated. But it would not be practical to attempt to catalog the enthalpy change for every possible chemical reaction. Instead, values for the heat of formation of compounds are tabulated. As long as these values are available for all of the substances involved, they can be used to calculate the enthalpy change for any reaction. The enthalpy change can also be related to the stoichiometry of a reaction, and such calculations can be used to find the amount of fuel needed to produce a given amount of energy.

KEY TERMS boundary (9.3)

formation reaction (9.5)

primary battery (9.8)

British thermal unit (Btu) (9.1)

heat (9.2)

PV-work (9.2)

calorie (cal) (9.2)

heat capacity (9.4)

secondary battery (9.8)

calorimetry (9.4)

heat of formation (9.5)

specific heat (9.4)

chemical energy (9.2)

heat of reaction (9.5)

state function (9.6)

endothermic (9.5)

Hess’s law (9.6)

surroundings (9.3)

energy density (9.7)

internal energy (9.2)

system (9.3)

energy economy (9.1)

joule (J) (9.2)

thermochemical equation (9.5)

enthalpy (9.5)

kinetic energy (9.2)

thermochemistry (9.1)

enthalpy diagram (9.6)

molar heat capacity (9.4)

universe (9.3)

exothermic (9.5)

oxidation–reduction reaction (9.8)

work (9.2)

first law of thermodynamics (9.3)

potential energy (9.2)



INSIGHT INTO Energy Use and the World Economy 9.1 List reasons why there might be a connection between the amount of energy used by a country and its economic development. 9.2 What are the advantages of electricity as a type of energy that make it worth generating despite the sizable losses that occur during this process? 9.3 What differences (if any) would you expect to find between the energy use patterns of the United States shown in Figures 9.2 and 9.3 and use in (a) European countries and (b) developing countries? 9.4 The total energy supply for the United States in the year 2007, as shown in Figure 9.2, was 106.96 × 10 15 Btu. (a) What percentage of that energy was obtained from imports? (b) What percentage was used to generate electricity? 9.5 Using the value you calculate from (a) in the previous problem and the fact that an average barrel of petroleum has an energy equivalent of 5.85 million Btu, how many barrels of oil, on average, must be imported each day?

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Defining Energy 9.6 Distinguish between kinetic and potential energy. 9.7 Define the term internal energy. 9.8

■ How fast (in meters per second) must an iron ball with a mass of 56.6 g be traveling in order to have a kinetic energy of 15.75 J? The density of iron is 7.87 g/cm3.

9.9 What is the kinetic energy of a single molecule of oxygen if it is traveling at 1.5 × 103 m/s? 9.10

The kinetic energy of molecules is often used to induce chemical reactions. The bond energy in an O2 molecule is 8.22 × 10–19 J. Can an O2 molecule traveling at 780 m/s provide enough energy to break the O"O bond? What is the minimum velocity of an O2 molecule that would give a kinetic energy capable of breaking the bond if it is converted with 100% efficiency? ■

9.11 Analyze the units of the quantity (pressure × volume) and show that they are energy units, consistent with the idea of PV-work. 9.12

■

How many kilojoules are equal to 3.27 L atm of work?

9.13 Define the term hydrocarbon.

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9.14 What are the products of the complete combustion of a hydrocarbon?

How does a hot day at the beach provide evidence for your answer?

9.15 Carry out the following conversions of energy units: (a) 14.3 Btu into calories, (b) 1.4 × 105 cal into joules, (c) 31.6 mJ into Btu

9.31 A metal radiator is made from 26.0 kg of iron. The specific heat of iron is 0.449 J/g °C. How much heat must be supplied to the radiator to raise its temperature from 25.0 to 55.0°C?

9.16 According to Figure 9.2, the total energy supply in the United States in 2007 was 106.96 × 1015 Btu. Express this value in joules and in calories. Energy Transformation and Conservation of Energy 9.17 If a machine does 4.8 × 103 kJ of work after an input of 7.31 × 104 kJ of heat, what is the change in internal energy for the machine? 9.18

■ Calculate (a) q when a system does 54 J of work and its energy decreases by 72 J and (b) DE for a gas that releases 38 J of heat and has 102 J of work done on it.

9.19 If the algebraic sign of DE is negative, in which direction has energy flowed? 9.20 State the first law of thermodynamics briefly in your own words.

9.32 The material typically used to heat metal radiators is water. If a boiler generates water at 79.5°C, what mass of water was needed to provide the heat required in the previous problem? Water has a specific heat of 4.184 J/g °C. 9.33 Copper wires used to transport electrical current heat up because of the resistance in the wire. If a 140-g wire gains 280 J of heat, what is the change in temperature in the wire? Copper has a specific heat of 0.384 J/g °C. 9.34 A copper nail and an iron nail of the same mass and initially at the same room temperature are both put into a vessel containing boiling water. Which one would you expect to reach 100°C first? Why? 9.35

9.21 Which type of energy, heat or work, is “valued” more by society? What evidence supports your judgment? 9.22 PV-work occurs when volume changes and pressure remains constant. If volume is held constant, can PV-work be done? What happens to Equation 9.2 when volume is held constant?

9.36 Define the term calibration. 9.37

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9.38

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9.23 Which system does more work: (a) DE = –436 J, q = 400 J; or (b) DE = 317 J, q = 347 J? 9.24 In which case is heat added to the system: (a) DE = –43 J, w = 40 J; or (b) DE = 31 J, w = 34 J? 9.25 Figure 9.5 shows projections for improved efficiency in many important technologies over the next two decades. Will these efficiency gains necessarily lead to reduced household energy demand? What factors might apply pressure for continued increases in household energy use? 9.26 Gas furnaces have achieved impressive efficiency levels largely through the addition of a second heat exchanger that condenses water vapor that would otherwise escape out the exhaust system attached to the furnace. How does this process improve efficiency? 9.27 When energy conservation programs are promoted, they sometimes include strategies such as turning out lights. Use the web to determine the average percentage of U.S. household energy usage that is attributable to lighting and use the information to comment on how effective the “turn out the lights” strategy might be. 9.28

■ When an electrical appliance whose power usage is X watts is run for Y seconds, it uses X × Y joules of energy. The energy unit used by electrical utilities in their monthly bills is the kilowatt-hour (kWh, that is, 1 kilowatt used for 1 hour). How many joules are there in a kilowatt-hour? If electricity costs $.09 per kilowatt-hour, how much does it cost per megajoule?

Heat Capacity and Calorimetry 9.29 Define the term calorimetry. 9.30 For the example of shallow water and sandy beaches, which material has a larger heat capacity or specific heat?

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■ A piece of titanium metal with a mass of 20.8 g is heated in boiling water to 99.5°C and then dropped into a coffee cup calorimeter containing 75.0 g of water at 21.7°C. When thermal equilibrium is reached, the final temperature is 24.3°C. Calculate the specific heat capacity of titanium.

A calorimeter contained 75.0 g of water at 16.95°C. A 93.3-g sample of iron at 65.58°C was placed in it, giving a final temperature of 19.68°C for the system. Calculate the heat capacity of the calorimeter. Specific heats are 4.184 J/g/°C for H2O and 0.444 J/g/°C for Fe. The energy densities of various types of coal are listed below. Anthracite 35 kJ/g

Subbituminous 31 kJ/g

Bituminous 28 kJ/g

Lignite 26 kJ/g

An unknown sample of one of these coals is burned in an apparatus with a calorimeter constant of 1.3 kJ/°C. When a 0.367-g sample is used, the temperature change is 8.75°C. Which type of coal is the sample? 9.39

■ How much thermal energy is required to heat all of the water in a swimming pool by 1°C if the dimensions are 4 ft deep by 20 ft wide by 75 ft long? Report your result in megajoules.

9.40 How does the specific heat of water explain why cities on the coast of large bodies of water tend to be cooler in the summer than cities several miles inland? Enthalpy 9.41 Under what conditions does the enthalpy change equal the heat of a process? 9.42 Why is enthalpy generally more useful than internal energy in the thermodynamics of real world systems? 9.43 Define the terms exothermic and endothermic. 9.44 List at least two phase changes that are exothermic processes. 9.45 What happens to the temperature of a material as it undergoes an endothermic phase change? If heat is added, how can the temperature behave in this manner?


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9.46 The heat of fusion of pure silicon is 43.4 kJ/mol. How much energy would be needed to melt a 5.24-g sample of silicon at its melting point of 1693 K?

9.58

5 (a) C2H2(g) + — O2(g) : 2 CO2(g) + H2O(𝓵) 2

9.47 If 14.8 kJ of heat is given off when 1.6 g of HCl condenses from vapor to liquid, what is DHcond for this substance?

(b) PCl3(g) + Cl2(g) : PCl5(g)

9.48 Calculate the energy required to convert 1.70 g of ice originally at –12.0°C into steam at 105°C. 9.49 DHvap = 31.3 kJ/mol for acetone. If 1.40 kg of water were vaporized to steam in a boiler, how much acetone (in kg) would need to be vaporized to use the same amount of heat? 9.50

■ When

a 13.0-g sample of NaOH(s) dissolves in 400.0 mL water in a coffee cup calorimeter, the temperature of the water changes from 22.6°C to 30.7°C. Assuming that the specific heat capacity of the solution is the same as for water, calculate (a) the heat transfer from system to surroundings and (b) DH for the reaction NaOH(s) : Na+(aq) + OH–(aq)

9.51 Define the term formation reaction. 9.52 Write the formation reaction for each of the following substances: (a) CH 4 (g), (b) C 3 H 8 (ℓ), (c) HCl(g), (d) C6H12O6(s), (e) NaF(s) 9.53 Explain why each of the following chemical equations is not a correct formation reaction: (a) 4 Al(s) + 3 O2(g) : 2 Al2O3(s)

(c) C2H4(g) + H2O(g) : C2H5OH(g) (d) Fe2O3(s) + 2 Al(s) : Al2O3(s) + 2 Fe(ℓ) 9.59 The heat of combustion of butane is –2877 kJ/mol. Use this value to find the heat of formation of butane. (You may also need to use additional thermochemical data found in Appendix E.) 9.60 When a chemical bond breaks, is energy absorbed or released? 9.61 When a reaction is exothermic, is the sum of bond energies of products or of reactants greater? Energy and Stoichiometry 9.62 For the reaction C2H2(g) + 2 H2(g) : C2H6, DH° = –136 kJ, what are the ratios that can be defined between moles of substances and energy? 9.63 For the reaction N2(g) + O2(g) : 2 NO(g), DH° = 180.5 kJ, how much energy is needed to generate 35 moles of NO(g)? 9.64

3 (b) N2(g) + — H2(g) : NH3(g) 2 (c) 2 Na(s) + O(g) : Na2O(s) Hess’s Law and Heats of Reaction 9.54

9.55

Which of the following are state functions? (a) the volume of a balloon, (b) the time it takes to drive from your home to your college or university, (c) the temperature of the water in a coffee cup, (d) the potential energy of a ball held in your hand

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Nitroglycerine, C 3H 5(NO 3) 3(,), is an explosive most often used in mine or quarry blasting. It is a powerful explosive because four gases (N 2, O 2, CO 2, and steam) are formed when nitroglycerine is detonated. In addition, 6.26 kJ of heat is given off per gram of nitroglycerine detonated. (a) Write a balanced thermochemical equation for the reaction. (b) What is DH when 4.65 mol of products is formed? ■

9.65 Silane, SiH4, burns according to the reaction, SiH4 + 2 O2 : SiO2 + 2 H2O, with DH ° = –1429 kJ. How much energy is released if 15.7 g of silane is burned?

■ Using these reactions, find the standard enthalpy change for the formation of 1 mol PbO(s) from lead metal and oxygen gas.

9.66 Sulfur trioxide can be removed from the exhaust gases of power plants by reaction with lime according to the equation, CaO(s) + SO3(g) : CaSO4(s), with DH ° = –886 kJ. If 240 kg of SO3 is to be removed, how much heat is released?

PbO(s) + C(graphite) : Pb(s) + CO(g)

9.67

2 C(graphite) + O2(g) : 2 CO(g)

DH° = 106.8 kJ DH°= –221.0 kJ

If 250. g of lead reacts with oxygen to form lead(II) oxide, what quantity of thermal energy (in kJ) is absorbed or evolved? 9.56 The phase change between graphite and diamond is difficult to observe directly. Both substances can be burned, however. From these equations, calculate DH° for the conversion of diamond into graphite. C(s, graphite) + O2(g) : CO2(g)

DH° = –393.51 kJ

C(s, diamond) + O2(g) : CO2(g)

DH° = –395.94 kJ

9.57 Hydrogen gas will react with either acetylene or ethylene gas. The thermochemical equations for these reactions are provided below. Write the thermochemical equation for the conversion of acetylene into ethylene by hydrogen gas.

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■ Using heats of formation tabulated in Appendix E, calculate the heats of reaction for the following:

C2H2(g) + 2 H2(g) : C2H6

DH° = –311 kJ

C2H4(g) + H2(g) : C2H6

DH° = –136 kJ


■ Reactions of hydrocarbons are often studied in the petroleum industry. One such reaction is 2 C 3H 8(g) : C6H6(ℓ) + 5 H2(g), with DH ° = 698 kJ. If 35 L of propane at 25°C and 0.97 atm is to be reacted, how much heat must be supplied?

9.68 In principle, ozone could be consumed in a reaction with lead and carbon with the thermochemical equation, Pb(s) + C(s) + O3(g) : PbCO3(s) DH ° = –841.0 kJ. How much energy would be released if 110 g of ozone reacts with excess lead and carbon? 9.69 When 0.0157 g of a compound with a heat of combustion of –37.6 kJ/mole is burned in a calorimeter, 18.5 J of heat is released. What is the molar mass of the compound? 9.70 Define the term energy density. 9.71 Why is energy density so important in the transportation of fuels? 9.72 What are some features of petroleum that make it such an attractive fuel?

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INSIGHT INTO Batteries 9.73 What type of chemical reaction takes place in a battery? 9.74 A small AAA battery and a much larger D battery both supply the same voltage. How does the underlying chemistry of the alkaline battery explain this? What factors would lead an engineer to specify D batteries in the design of a particular device rather than the smaller AAA-cells?

9.79 A student performing a calorimetry experiment combined 100.0 mL of 0.50 M HCl and 100.0 mL of 0.50 M NaOH in a Styrofoam™ cup calorimeter. Both solutions were initially at 20.0°C, but when the two were mixed, the temperature rose to 23.2°C. (a) Suppose the experiment is repeated in the same calorimeter but this time using 200 mL of 0.50 M HCl and 200.0 mL of 0.50 M NaOH. Will the DT observed be greater than, less than, or equal to that in the first experiment, and why?

9.75 Explain the meaning of the terms primary and secondary in the context of batteries. Give examples of both types of batteries. 9.76 Consumers face a variety of choices in selecting batteries. A typical AA alkaline battery might have a rated lifetime of about 2800 mA hours and cost around $0.50. New high-capacity alkaline cells can last for 3100 mA hours but typically cost about $1.00 each. Rechargeable nickel metal hydride (NiMH) cells offer yet another option, with a lifetime of 2100 mA hours (per charge cycle) and cost around $3.50 each. What might lead you to choose each of these options? What type of applications do you think would be best suited for each of these battery types?

(b) Suppose that the experiment is repeated once again in the same calorimeter, this time using 100 mL of 1.00 M HCl and 100.0 mL of 1.00 M NaOH. Will the DT observed be greater than, less than, or equal to that in the first experiment, and why? 9.80 The specific heat of gold is 0.13 J g–1 K–1, and that of copper is 0.39 J g–1 K–1. Suppose that we heat both a 25-g sample of gold and a 25-g sample of copper to 80°C and then drop each into identical beakers containing 100 mL of cold water at 10°C. When each beaker reaches thermal equilibrium, which of the following will be true, and why? You should not need to do any calculations here.

Additional Problems 9.77

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The figure below shows a “self-cooling” beverage can.

(i) Both beakers will be at the same temperature. (ii) The beaker with the copper sample in it will be at a higher temperature.

Fill with 100 g H2O

(iii) The beaker with the gold sample in it will be at a higher temperature.

200 g beverage

9. 81

55 g Na2CO3

The can is equipped with an outer jacket containing sodium carbonate (Na2CO3), which dissolves in water rapidly and endothermically. Na2CO3(s) : 2 Na+ (aq) + CO32–(aq)

9.78

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You make some iced tea by dropping 134 g of ice into 500.0 mL of warm tea in an insulated pitcher. If the tea is initially at 20.0°C and the ice cubes are initially at 0.0°C, how many grams of ice will still be present when the contents of the pitcher reach a final temperature? The tea is mostly water, so assume that it has the same density (1.0 g/mL), molar mass, heat capacity (75.3 J K–1 mol–1), and heat of fusion (6.0 kJ/mol) as pure water. The heat capacity of ice is 37.7 J K–1 mol–1. ■


9.82 Chemical engineers often must include ways to dissipate energy in the form of heat in their designs. What does this fact say about the enthalpy change of chemical reactions that are being used? 9.83

DH ° = 67.7 kJ

The user adds water to the outer jacket, and the heat absorbed in the chemical reaction chills the drink. The can contains 200 g of drink, the jacket contains 55 g of Na2CO3, and 100 g of water is to be added. If the initial temperatures of the can and the water are 32°C on a summer day, what is the coldest temperature that the drink can reach? The can itself has a heat capacity of 40 J/°C. Assume that the Na2CO3 solution and the drink both have the same heat capacity as pure water, 4.184 J g–1 °C–1. (HINT: Treat this like a calorimetry problem.)

■ What will be the final temperature of a mixture made from equal masses of the following: water at 25.0°C, ethanol at 35.5°C, and iron at 95°C?

A sample of natural gas is 80.0% CH4 and 20.0% C2H6 by mass. What is the heat from the combustion of 1.00 g of this mixture? Assume the products are CO2(g) and H2O(,).

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9.84 Many engineering designs must incorporate ways to dissipate energy in the form of heat. Water evaporators are common for this task. (a) What property of water makes it a good material for evaporators? (b) If an application could not use water, but instead was forced to use a material with a value for the property in (a) that was one half that of water, what changes would need to be made in the design? 9. 85

You want to heat the air in your house with natural gas (CH4). Assume your house has 275 m2 (about 2800 ft2) of floor area and that the ceilings are 2.50 m from the floors. The air in the house has a molar heat capacity of 29.1 J/mol/K. (The number of moles of air in the house v be found by assuming that the average molar mass of air is 28.9 g/mol and that the density of air at these temperatures is 1.22 g/L.) What mass of methane do you have to burn to heat the air from 15.0°C to 22.0°C? ■

9.86 The curing of concrete liberates energy as heat. (a) What does this observation suggest happens in terms of


Copyright 2009 Cengage Learning, Inc. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part.

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9. 87

chemical bonds as concrete cures? (b) What type of strategies might a civil engineer employ in designs and construction specifications to mitigate heat expansion of concrete as it cures?

observation to infer something about the specific heat of the materials in the baking sheets? (b) What is the mathematical reasoning (equation) that you need to support your conclusion?

■ An engineer is designing a product in which a copper wire will carry large amounts of electricity. The resistive heating of a 65-g copper wire is expected to add 580 J of heat energy during a 10-minute operating cycle. The specific heat of copper is 0.385 J g–1 °C–1, the density is 8.94 g/cm3, and the coefficient of thermal expansion is 16.6 μm m–1 K–1. (a) What is the temperature increase of the wire? (b) What is the initial length of the wire, assuming it is a cylinder and its radius is 0.080 cm? (c) By what percentage does the length increase because of the temperature increase? (d) Do you think the engineer should be considering this expansion in the design?

9.93 Silicon nitride, Si3N4, has physical, chemical, and mechanical properties that make it a useful industrial material. For a particular engineering project, it is crucial that you know the heat of formation of this substance. A clever experiment allows direct determination of the DH ° of the following reaction:

9.88 Price spikes in gasoline in 2008 led to renewed interest in coal gasification projects, in which coal is converted to gasoline. Looking at the relative energy density of gasoline and coal in Table 9.4, which is more likely required in an engineering design for this project: the ability to input heat or the need to dissipate heat? Explain your reasoning.

FOCUS ON PROBLEM SOLVING EXERCISES 9.89 Suppose that the working fluid inside an industrial refrigerator absorbs 680 J of energy for every gram of material that vaporizes in the evaporator. The refrigerator unit uses this energy flow as part of a cyclic system to keep foods cold. A new pallet of fruit with a mass of 500 kg is placed in the refrigerator. Assume that the specific heat of the fruit is the same as that of pure water because the fruit is mostly water. Describe how you would determine the mass of the working fluid that would have to be evaporated to lower the temperature of the fruit by 15°C. List any information you would have to measure or look up. 9.90 Hydrogen combines with oxygen in fuel cells according to the thermochemical equation 2 H2(g) + O2(g) : 2 H2O(g)

DH ° = – 571.7 kJ

Suppose that you are working with a firm that is using hydrogen fuel cells to power satellites. The satellite requires 4.0 × 105 kJ of energy during its useful lifetime to stabilize its orbit. Describe how you would determine the mass of hydrogen you would need in your fuel cells for this particular satellite. 9.91 The chemical reaction, BBr3(g) + BCl3(g) : BBr2Cl(g) + BCl2Br(g), has an enthalpy change very close to zero. Using Lewis structures of the molecules, all of which have a central boron atom, provide a molecular level description of why DH ° for this reaction might be very small. 9.92 Two baking sheets are made of different metals. You purchase both and bake a dozen cookies on each sheet at the same time in your oven. You observe that after 9 minutes, the cookies on one sheet are slightly burned on the bottom, whereas those on the other sheet are fine. (You are curious and you vary the conditions so you know the result is not caused by the oven.) (a) How can you use this

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3 CO2(g) + Si3N4(s) : 3 SiO2(s) + 2 N2(g) + 3 C(s) Based on the fact that you know the enthalpy change in this reaction, state what additional data might be looked up or measured to determine DHf° for silicon nitride. 9.94 A runner generates 418 kJ of energy per kilometer from the cellular oxidation of food. The runner’s body must dissipate this heat or the body will overheat. Suppose that sweat evaporation is the only important cooling mechanism. If you estimate the enthalpy of evaporation of water as 44 kJ/mol and assume that sweat can be treated as water, describe how you would estimate the volume of sweat that would have to be evaporated if the runner runs a 10-km race. 9.95 One reason why the energy density of a fuel is important is that to move a vehicle one must also move its unburned fuel. Octane is a major component of gasoline. It burns according to the reaction, 2 C8H18(ℓ) + 25 O2(g) : 16 CO2(g) + 18 H2O(g) DH ° = –1.10 × 104 kJ Starting from this thermochemical equation, describe how you would determine the energy density, in kJ/g, for octane. Be sure to indicate what you would need to calculate or look up to complete this problem. 9.96 An engineer is using sodium metal as a cooling agent in a design because it has useful thermal properties. Looking up the heat capacity, the engineer finds a value of 28.2 J mol–1 °C–1. Carelessly, he wrote this number down without units. As a result, it was later taken as specific heat. (a) What would be the difference between these two values? (b) Would the engineer overestimate the ability of sodium to remove heat from the system or underestimate it because of this error? Be sure to explain your reasoning. 9.97 In passive solar heating, the goal is to absorb heat from the Sun during the day and release it during the night. Which material would be better for this application: one with a high heat capacity or one with a low heat capacity? Explain. 9.98 A 1.0-kg sample of stainless steel is heated to 400°C. Suppose that you drop this hot sample into an insulated bucket that contains water at some known initial temperature. Assuming that there is no difficulty in transferring heat from the steel to the water, describe how you can determine the maximum mass of water that could be boiled with only the heat given off by this sample of steel. Be sure to list any quantities you would need to look up to solve this problem.

■


Cumulative Problems 9.99

■ At the beginning of 2008, the United States had 21.3 billion barrels of proven oil reserves. One barrel of oil can produce about 19.5 gallons of gasoline. Assume that the gallon of gasoline is pure octane, with a density of 0.692 g/mL. If all 21.3 billion barrels of oil were converted to gasoline and burned, how much energy would be released?

9.100

■

9.101

■

■

For a car weighing 980 kg, how much work must be done to move the car 24 miles? Ignore factors such as frictional loss and assume an average acceleration of 2.3 m/s2. Suppose that the car in the previous problem has a fuel efficiency of 24 mpg. How much energy is released


in burning a gallon of gasoline (assuming that all of it is octane)? Based on this calculation and the work required to move the car those 24 miles, what percentage of the energy released in the combustion is wasted (doesn’t directly contribute to the work of moving the car)? How do the assumptions you make in carrying out these calculations affect the value you obtain? 9.102 Suppose that there is 2.43 mol of nitrogen gas in an insulated, sealed 31.7-L container initially at 285 K. The specific heat of nitrogen gas is 1.04 kJ kg –1 K –1 (note units). If a 5.44-kg block of iron at 755 K is placed in this container and it is sealed again (with no loss of nitrogen), what is the final pressure of the nitrogen gas?


317

10 OUTLINE

Entropy and the Second Law of Thermodynamics

10.1 INSIGHT INTO Recycling of Plastics 10.2 Spontaneity 10.3 Entropy 10.4 The Second Law of Thermodynamics 10.5 The Third Law of Thermodynamics 10.6 Gibbs Free Energy 10.7 Free Energy and Chemical Reactions 10.8 INSIGHT INTO The Economics of Recycling

Curbside recycling programs often collect “comingled” materials, as seen here at a Milwaukee site. Plastics, which make up about 85% of this pile, must be separated and sorted for recycling. Thomas A. Holme

I


318

n our discussions of chemical bonding, we introduced the idea that bonds form because doing so reduces the overall energy of the collection of atoms involved. We’ve seen many examples of chemical reactions, such as combustion and explosions, which also reduce the overall energy of the atoms and molecules involved. But if you try for just a moment, you should also be able to think of many common chemical and physical processes in which the energy of the system clearly increases. Ice cubes melt. The batteries in your laptop or cell phone are recharged. At least some endothermic chemical reactions occur regularly. As each of these cases shows, the energy of a system does not always decrease, despite our intuitive sense of a preference for minimizing energy. So how can we understand and predict which changes nature will actually favor? We will need to extend our understanding by introducing the second law of thermodynamics and exploring its ramifications. Although applications abound in virtually all fields of science and engineering, the impact of thermodynamics on our understanding of chemical reactions has been especially profound. We’ll explore the implications of the second law by looking into the recycling of plastics.

Chapter Objectives After mastering this chapter, you should be able to ❚ describe the scientific and economic obstacles to more widespread recycling of plastics. ❚ explain the concept of entropy in your own words. ❚ deduce the sign of DS for many chemical reactions by examining the physical state of the reactants and products. ❚ state the second law of thermodynamics in words and equations and use it to predict spontaneity. ❚ state the third law of thermodynamics. ❚ use tabulated data to calculate the entropy change in a chemical reaction. ❚ derive the relationship between the free energy change of a system and the entropy change of the universe. ❚ use tabulated data to calculate the free energy change in a chemical reaction. ❚ explain the role of temperature in determining whether a reaction is spontaneous. ❚ use tabulated data to determine the temperature range for which a reaction will be spontaneous.

INSIGHT INTO

10.1

Recycling of Plastics

Standard plastic soft drink bottles are made of poly(ethylene terephthalate), or PET. In the industrial-scale synthesis of PET, the usual starting materials are dimethyl terephthalate and ethylene glycol (Figure 10.1). These compounds react to form bis(2-hydroxyethyl) terephthalate (BHET) and methanol. The methanol boils off at the reaction temperature (typically around 210°C), leaving fairly pure BHET. Then, the BHET is heated further to around 270°C, where it undergoes a condensation reaction to form PET polymer. Ethylene glycol is a byproduct in this second step and can thus be reused within the plant to produce more BHET. The resulting polymer can be melted, blown, and molded into bottles of the desired shape. These bottles are then filled, capped, shipped, and sold. You pick up your O

O

H3C

O

O

O

HO ⫹

Condensation reactions were introduced in Section 8.6.

O ⫹

2

OH

CH3

O

O

HO Dimethyl terephthalate (DMT)

O

Ethylene glycol

O

OH bis-(2-Hydroxyethyl) terephthalate (BHET)

O

O

O

[ O

O

HO

n

OH

O

]n

OH bis-(2-Hydroxyethyl) terephthalate (BHET)

Methanol

HO ⫹

n

2 CH3OH

Poly(ethylene terephthalate) (PET)

Ethylene glycol

Figure 10.1 ❚ Steps in the industrial synthesis of PET are illustrated. Typical values of n in the polymer formula are 130–150, giving a molar mass of around 25,000 for the polymer.

10.1 Recycling of Plastics

319

Some of the crushing and sorting is now done automatically in “reverse vending machines” designed to collect bottles for recycling.

Many manufacturers have introduced bottles with new rounder designs that allow the use of thinner plastic.

soda, perhaps from a vending machine on your way to class. Once the bottle is empty, you toss it into the recycling bin and feel good that you’ve done your part to help protect the environment. Chances are you may never have thought much about what happens to the bottle from there. Typically, the contents of the recycling bin are sold to a reclaimer—a business specializing in processing plastics. Usually, the material from the bin must be sorted into different types of plastics, and any other materials that may have been thrown into the bin are discarded. Some of this sorting is done by hand, and some takes advantage of differences in density among the various polymers that might be present. The plastics are then crushed to reduce their volume before being shipped for further processing. The next step is called reclamation, in which the sorted and compressed plastics are processed into a useable form. In most reclamation processes, the plastic is first chopped into small, uniform-sized flakes. These flakes are washed and dried, then melted and extruded into spaghetti-like strands. These are then cut into smaller pellets, which are sold to manufacturers for use in new products. The most important uses for recycled PET include fiberfill for sleeping bags and coats, fleece fabrics for outdoor wear, carpeting, and industrial strapping. You may have noticed that one thing does not appear on that list of uses: new drink bottles. Although such bottles are the dominant source of PET for recycling, only very limited amounts of recycled PET are used to make new bottles. Thus the recycling of PET is far from being a “closed loop” process; large amounts of virgin plastic continue to be used in bottling despite increased collection of used bottles at the consumer level. Why is this? The simplest and shortest answer is economics: bottles can be made from virgin plastic at a lower overall cost. Several factors contribute to this. In many cases, there are legal restrictions on the use of recycled materials for food and beverage containers, due to concerns over possible contamination. Satisfying these regulations adds cost to the overall equation. Degradation of the plastic during repeated recycling processes is another concern. The average chain length of the polymer molecules tends to be somewhat lower after recycling. So if bottles were made from 100% recycled PET, they might have to be thicker and heavier. Although progress is being made in increasing the recycled content of drink bottles, most U.S. bottles still contain at least 90% virgin plastic. One possible way to achieve a closed loop in which plastic bottles could be recycled back into plastic bottles might be to convert the polymer into monomers and then repolymerize the monomers to produce new plastic. Under what circumstances might such a scheme be feasible? Before we can explore that type of question, we will first need to learn more about thermodynamics.

10.2

Spontaneity

Nature’s Arrow The idea of time travel drives the plot in many science fiction stories. The prospect of moving forward or backward in time and existing in some other era appeals to our imagination in a way that provides fertile ground for authors. But our actual experience is that time marches inexorably from the past toward the future and that this direction is not reversible. In a sense, time is an arrow that points in the direction in which nature is headed. We’ve seen that large hydrocarbon molecules, such as those in gasoline, can react readily with oxygen to produce carbon dioxide and water. But your experience also tells you that the reverse reaction doesn’t happen; water vapor and carbon dioxide are always present in the air, but they never react to produce gasoline. Nature clearly “knows” the correct direction for this process. This sense of the direction of life and our experience of the universe is an important intuition to carry into this chapter. But what gives nature this direction? And how can we convert our intuition into a useful quantitative model for predicting which chemical reactions will actually occur? We’ll try to answer these questions by imparting a bit of mathematical rigor to our observations. 320

Chapter 10 Entropy and the Second Law of Thermodynamics

Spontaneous Processes A more formal way of expressing the directionality of nature is to note that our intuition is predicated on the fact that some things “just happen,” but others do not. Some processes occur without any outside intervention, and we say that such a process is spontaneous. From a thermodynamic perspective, then, a spontaneous process is one that takes place without continuous intervention. The distinction between spontaneous and nonspontaneous reactions may seem obvious, but we’ll see that it is not always so. Students often misinterpret the word spontaneous as indicating that a process or reactions will take place quickly. But note that our actual definition does not refer to the speed of the process at all. Some spontaneous processes are very fast, but others occur only on extremely long timescales. We understand that the chemical compounds in some waste materials, like paper, may spontaneously react to decay over time. (This process can be more complicated than a simple chemical reaction, though, because of the involvement of bacteria.) But some spontaneous reactions are so slow that we have a hard time observing them at all. The combustion of diamond is thermodynamically spontaneous, yet we think of diamonds as lasting forever. Other reactions occur quickly once they start, but they don’t just start on their own. Gasoline, for example, can sit more or less indefinitely in a can in the garage, in contact with oxygen in the air. Nonetheless, no reaction is observed. Yet, upon being mixed with air in the cylinder of your car and ignited by the spark plug, the reaction proceeds until virtually all the gasoline is burned. Is this reaction spontaneous? The answer is yes. Even though the reaction needs a flame or spark to initiate it, once it begins, the reaction continues without any further intervention. This example emphasizes the importance of the phrase “continuous intervention” in our definition. A useful analogy is that of a rock perched precariously on a cliff. If it is nudged over the edge, it proceeds to the bottom. It does not stop midway down, unless, of course, it’s a prop in a Roadrunner cartoon! The reactions used to produce many polymers behave much like the combustion of gasoline. Once initiated, the reaction is usually spontaneous and can proceed without further intervention. The production of poly(methyl methacrylate)—Plexiglas, or PMMA—is a good example. H n C

H

CH3

[

C C

H

O

CH3 C C

We mentioned PMMA in Section 7.1 as having been used as one of the first bone cements.

C O

H

O

[n

Some spontaneous processes take place over geological time scales—the formation of petroleum used for plastics feedstocks, for example.

O CH3

CH3

Methyl methacrylate monomer

Poly(methyl methacrylate)

This reaction occurs via a free radical process, like that described in Section 2.8 for polyethylene. A small trace of an initiator is needed to start the reaction, and then it proceeds until virtually all of the available monomer has been converted into polymer. But suppose that we wanted to convert the polymer back into monomer. In that case, the necessary reaction is the reverse of the polymerization, and it is not a thermodynamically spontaneous process at ordinary temperatures. We could still drive the reaction backward to produce methyl methacrylate monomer. But we would need to maintain a high temperature, providing enough energy to allow the molecules to go against nature’s preferred direction. So what is the role of energy in the directionality of nature?

Enthalpy and Spontaneity Recall from Chapter 9 that the enthalpy change in a chemical reaction is equal to the heat flow at constant pressure. DH 5 qp 10.2

Spontaneity

321

When DH is negative, the reaction is exothermic, whereas a positive value of DH points to an endothermic reaction. What can we say about a reaction’s spontaneity based on its enthalpy change? If we were to stop and list spontaneous processes that we observe around us and then determine whether those processes are exothermic or endothermic, chances are that a majority would be exothermic. This implies that there is some relationship between enthalpy and spontaneity. The relationship is not exclusive, however. If you think for a moment you should be able to point out some endothermic reactions that obviously occur spontaneously. The melting of an ice cube at room temperature is one simple example. So at this point we might conclude that exothermic reactions seem to be preferred in some way. But clearly there must be things other than energy or enthalpy at work in determining whether or not a process is spontaneous. To develop a way to predict the spontaneity of a reaction, we must first introduce an additional thermodynamic state function—entropy.

10.3 A state function does not depend on the system’s history. So there can be no change in any state function for a process where the initial and final states are the same.

Entropy

As we have just seen, the flow of energy as heat does not indicate whether or not a process will occur spontaneously. So we must also consider another thermodynamic state function, called entropy. Historically, entropy was first introduced in considering the efficiency of steam engines. Figure 10.2 illustrates the Carnot cycle, which uses a combination of adiabatic processes (in which no heat is exchanged) and isothermal, or constant temperature, processes. The Carnot cycle demonstrated that a previously unknown state function existed because the sum of q/T (heat divided by temperature) around the closed path is zero. This new state function was called entropy. We’ll soon see that the changes in the entropy of a system and its surroundings allow us to predict whether or not a process is spontaneous. What is entropy and how can it help us understand the production or recycling of polymers?

Probability and Spontaneous Change We can observe a pattern in many changes that occur in everyday life that are analogous to events at a molecular level. For a familiar example, let’s think about autumn. Although the falling leaves may be welcome as a sign of cooler temperatures, they also mean an added chore—raking the leaves into piles. Why can’t the leaves simply fall in a pile to begin with? Such an event goes against our intuition because it is so unlikely that we know we’ll Figure 10.2 ❚ In the Carnot cycle, an ideal gas undergoes a series of four processes. Two of these (labeled 1 and 3 in the figure) are isothermal, which means they occur at constant temperature. The other two steps (2 and 4) are adiabatic, meaning that q 5 0 for those parts of the cycle. Carnot showed that the sum of the quantity q/T for the entire cycle is equal to zero. Because the cycle begins and ends with the system in the same state, this means that there must be a state function equal to q/T. We call this state function entropy.

P

1

Isothermal expansion q/T > 0

Adiabatic 2 expansion q/T = 0

Adiabatic 4 compression q/T = 0 3 Isothermal contraction q/T < 0

V The sum of q/T around the cycle is zero, so there must be a state function that is given by this expression.

322


never see it happen. This macroscopic example with perhaps thousands of leaves serves as a reasonable analogy for molecular systems with Avogadro’s number of particles. Let’s look at the concept of mathematical probability to solidify our understanding. The example of leaves not falling in a pile, though perhaps obvious, is somewhat challenging to describe in mathematical terms. To establish a foundation in ideas of probability, let’s think instead about rolling dice. If you take just one die and roll it, what is the chance that the roll will be a four? With six possible outcomes the chance is one in six. For two dice, what is the chance the roll will be a pair of fours? This time the counting is a bit more involved, but we can quickly see from Figure 10.3 that the chances are 1 in 36. If a third die is added, the chances of rolling three fours in one throw are 1 in 216. We can see the relationship that is developing for rolling all fours. There is only one way to achieve it, and the probability of that outcome grows smaller according to the following relationship: 1 N, where N is the number of dice being thrown Probability 5 q—— r 6 We should note that this relationship applies for the case at hand, but it is not general. The factor of 1 in the numerator is present because we are looking for a single specific roll (of a number four) on each die and the six in the denominator is there because there are six possible rolls for each die. With this relationship, however, we could easily predict that the chances of rolling the same number with five dice in one roll are one in 1296. (Note that the chances of rolling a specified number on all five dice—say all fours—are 1 in 7776. But if we do not specify in advance which of the six possible numbers we want on all five dice, then there will be six possible outcomes instead of just one.) Our experience with rolling dice is that we expect to have some random assortment of numbers present when five dice are rolled. Why? There are very many ways to obtain a “random” roll. Such a roll occurs far more often precisely because it is more probable. Roll

# of ways

Roll

# of ways

1 2

6

3

5 4

4

Figure 10.3 ❚ The probability of rolling a given total value on a pair of dice depends on the number of different combinations that produce that total. The least likely rolls are 2 and 12, for example, because there is only one possible combination that gives each of those totals. The most likely total is seven because there are six different rolls that add up to that number. (Note that for rolls in which the two individual dice show different values, two possibilities exist. For a total roll of three, for example, the two combinations would be 1, 2 and 2, 1.)

5

3

2 1

10.3

Entropy

323

Our development of the mathematics of probability has two important features. First, it shows that to obtain the probability of a collection of events based on the probability of an individual event we must multiply. This observation becomes important when we consider just how many molecules are involved when we observe something in nature or in the laboratory. Second, the number of ways to make an ordered observation (like all dice turning up four) is smaller than the number of ways to make a more random observation (no particular pattern present in the dice). When we apply these observations to a collection of molecules with ,1023 particles present, the chances for highly specified arrangements become phenomenally small. To start addressing numbers with more chemical relevance, imagine rolling Avogadro’s num23 ber of dice. The probability of all of them coming up four is q—16 r6.02 3 10 . That number is unimaginably small. If we used all the zeros after the decimal point to replace all the letters in all the books on the planet, we would still have zeros left over!

Definition of Entropy For large numbers of particles, then, probability favors random arrangements. Using this insight, we can tentatively define entropy as a measure of the randomness or disorder of a system. However, we still have to establish a definition that can be used quantitatively and from a molecular perspective. To do this, we turn to a branch of physical chemistry called statistical mechanics, or statistical thermodynamics, where we find a subtle addition to the definition. The probability of events that must be counted is not the number of ways particles can be arranged physically but rather the number of ways in which particles can achieve the same energy. (These two probabilities are often correlated with one another.) If we recall the Maxwell-Boltzmann distribution of molecular speeds (see Section 5.6), we know that in any gas at room temperature, some particles must move slowly and others quite rapidly. We cannot, however, say precisely which particle is moving very fast or which particle is moving more slowly (Figure 10.4). There are a large number of different ways, with different particles assuming the various required speeds, that the sample can have the same total energy and hence the same temperature. In statistical mechanics, the way by which the collection of particles assumes a given energy is associated with a concept called a microstate. The number of microstates for a given energy is commonly designated by the uppercase Greek letter omega (V), and the entropy (S ) of a system is related to the number of microstates by the equation, S 5 kB ln V

(10.1)

Here kB is a numerical constant called the Boltzmann constant. It is not easy to have an intuition about the number of microstates of a system, so this equation is hard to use directly at this stage of your study of chemistry. We’ll soon see that we won’t need to use it. It is important, however, to realize that as a system becomes “more random,” the value of V will increase. So, the entropy of a system increases as the system moves toward more random distributions of the particles it contains because such randomness increases the number of microstates.

Judging Entropy Changes in Processes

The entropy of one mole of gas is generally very much greater than that of one mole of liquid or solid. 324

Although the concept of a microstate is abstract, we can still assert that certain types of changes will lead to increases in entropy (because there are more available microstates). Let’s see why this is so. First consider the melting of a solid to form a liquid. As a solid, the particles are held in place rigidly, so the number of ways they can have a specific energy is limited. When the liquid forms, the movement of particles relative to each other presents a much greater number of ways to achieve a specific energy, so the number of microstates increases and so does the entropy. Similar reasoning can be applied to boiling, when molecules originally confined near each other in a liquid become much more randomly distributed in the gas phase. The increase in random


Figure 10.4 ❚ The Maxwell-

The darker atom has a moderate velocity and the lighter atom has a high velocity. Both are part of the overall distribution.

Boltzmann distribution tells us the overall collection of molecular speeds but does not specify the speed of any individual particle. Energy exchange during molecular collisions can change the speed of individual molecules without disrupting the overall distribution.

Velocity

The darker atom now has a high velocity, but the overall distribution stays the same.

Velocity The lighter atom now has a low velocity, but the overall distribution stays the same.

molecular motion corresponds to more microstates, so entropy increases. Another possible way to increase the entropy of a system is to increase the number of particles present. Thus, a chemical reaction that generates two moles of gas where only one was present initially will increase the entropy. Entropy also varies with temperature. One way to think about this is to begin by considering a sample of molecules at some extremely low temperature. In such a sample, it would be very unlikely to have molecules moving at high speeds because they would account for too large a percentage of the available energy. So the speeds of individual molecules would be constrained by the low total energy available. If the system were heated to a higher temperature, though, then a few of the molecules could move at high speeds because there is more total energy available. We have only considered a small portion of the distribution of speeds of the molecules, but already we can see that a hotter system has more ways to distribute its energy. This type of reasoning extends to the whole distribution of speeds, and the important result is that heating a system increases its entropy. What are the implications of entropy for polymer synthesis or recycling? When a polymer is formed, a large number of monomers are converted into a single giant molecule. In most cases, this will lead to a decrease in the entropy of the system because there are more possible ways to arrange the unreacted monomers. (Note that in many polymerization reactions, other small molecules, such as water, may be formed as byproducts. In such cases, the sign of the entropy change may not be obvious.) The fact that polymerization reactions are still spontaneous under appropriate conditions tells us that entropy of the system alone is not the only important consideration. Other factors, such as energy, must favor the formation of the polymer. What about the role 10.3

Entropy

325

of entropy in recycling? As plastics are recycled, there is a possibility that the long polymer chains may be broken. From the viewpoint of entropy this should be a favorable process. Breaking the chains gives a smaller average molecular size, and the same system of atoms will have more microstates available if more individual molecules are present. Reducing the polymer chain length tends to weaken the mechanical properties of the plastic, though, because the shorter chains do not interact with one another as strongly. So entropy provides a challenge to the recycling process. To recycle polymers without a steady loss in the quality of the material, we need to overcome nature’s preference for increasing entropy. And as we’ll see by the end of this chapter, this leads very directly to real economic obstacles to recycling.

10.4

The Second Law of Thermodynamics

In Chapter 9, we stressed the importance of thermodynamics in terms of the way human society uses energy. At that time, we noticed that whenever there is an attempt to convert energy from one form into another, some energy is lost or wasted. In other words, not all of the energy potentially available is directed into the desired process. How does this fact arise from thermodynamics? Entropy provides the key to understanding that the loss of useful energy is inevitable.

The Second Law There are several equivalent ways to state the second law, but all lead to the same interpretations.

In considering the energy economy, we alluded to the second law in conjunction with the notion that it is impossible to convert heat completely to work. That is one way to express the second law of thermodynamics. Now let’s try to understand why this is true. First consider heat. Heat flows due to random collisions of molecules, and an increase in temperature increases the random motions of molecules. Work, by contrast, requires moving a mass some distance. To yield a net movement, there must be a direction associated with a motion, and that direction implies that there is an order to the motion. Converting heat into work, therefore, is a process that moves from random motions toward more ordered ones. We have just seen how this type of change goes against nature’s tendency to favor a more probable state (the more random one). How can we connect these ideas with entropy? To make this connection, we must be careful to realize that changes in the universe involve both the system and its surroundings. If we focus on the system alone, we cannot understand how order is created at all. Yet, the synthesis of polymers shows that it does happen, as do everyday situations such as the growth of plants, animals, and people. To express the second law of thermodynamics in terms of entropy, we must focus on the total change in entropy for the universe, DSu. DSu 5 DSsys 1 DSsurr

(10.2)

Because nature always tends to proceed toward a more probable state, we can assert an equivalent form of the second law of thermodynamics: In any spontaneous process, the total entropy change of the universe is positive, (DSu . 0). That this statement of the second law is equivalent to our original version is not at all obvious. But remember, energy that is not converted into work (a process that would decrease entropy) is transferred to the surroundings as heat. Thus the entropy of the surroundings increases, and the total entropy change in the system and surroundings is positive.

Implications and Applications The implications of this expression of the second law are far-reaching for calculating and predicting the outcome of chemical reactions and other processes we might wish to study. We can focus on these implications qualitatively first, by considering a 326


polymerization reaction from the perspective of thermodynamics. Later, we will develop a quantitative approach. Let’s return to the polymerization of methyl methacrylate to form PMMA. The monomer and polymer structures are shown on page 321. Looking at both structures, we see that most of the chemical bonds are unchanged during this reaction. The only exception is the C"C double bond in the monomer, which is converted into two C!C single bonds in the polymer. From our knowledge of bond energies, then, we can say that this reaction must be exothermic. Two C!C single bonds are stronger than a C"C double bond. Because the reaction converts a large number of monomer molecules into a single polymer molecule, we can also predict that the entropy change for the system must be negative. So why is the reaction spontaneous? The fact that the reaction is exothermic means that heat must be released from the system. That same heat must flow into the surroundings. This will lead to an increase in the entropy of the surroundings. As long as that increase in the entropy of the surroundings is larger than the decrease in the entropy of the system, the overall change in entropy for the entire process can still be positive. We can take this reasoning a little further to begin to understand the role of temperature in determining the spontaneity of a process. The surroundings will absorb an amount of heat equal to 2DH. But the surroundings represent a very large reservoir, so this heat will not produce a measurable temperature change. This means that the entropy change for the surroundings is given by DH DSsurr 5 2 —— T The entropy change for the system is just DS, and although we don’t know its value, we do know that it will be negative. The criterion for a spontaneous polymerization is DSu 5 DS 1 DSsurr . 0 This will be true as long as the absolute value of DSsurr is greater than that of DS. (Remember that DS is negative and DSsurr is positive.) The magnitudes of DS and DH are essentially independent of the temperature, but DSsurr will decrease as the temperature increases. So at some sufficiently high temperature, DSu will no longer be positive, and the reaction will cease to be spontaneous. This same argument points to a possible route to depolymerization, which might be useful in recycling. Suppose that we raise the temperature high enough that DSu for the polymerization reaction becomes negative. That must mean that DSu for the reverse reaction in which polymer is converted back into monomer must become positive. So if we heat the polymer above some threshold temperature, we should be able to regenerate methyl methacrylate monomer. When PMMA is heated above about 400°C, it is converted into monomer with a very high efficiency. This process, called thermolysis, is one example of what is often called advanced recycling or feedstock recycling. Because the recovered monomer can be purified by distillation or other means, it can then be repolymerized to produce PMMA that is indistinguishable from virgin material. Thermolysis is not practical for most plastics, though, because the monomers themselves often break down or undergo other undesirable reactions at the high temperatures that are required. In the particular case of PMMA, thermolysis is used mainly within manufacturing plants to reclaim scraps that are left behind in the production of items such as automobile taillight lenses.

10.5

Purification involves separating the monomer from other substances. Distillation is a common method for separation of chemical mixtures.

The Third Law of Thermodynamics

Thus far we have taken a purely qualitative approach to entropy changes and have not attempted to find numerical values for DS. To move toward a quantitative view, though, all we really need is to define some reference point with a fixed value of entropy. Then, as long as we can calculate entropy changes, we should be able to obtain values of interest. When we seek to calculate entropy changes for chemical reactions, the most 10.5


327

Table

❚

10.1

Standard molar entropies (S°) for selected substances. A much larger listing appears in Appendix E. Values for many compounds can also be found online in the NIST Chemistry WebBook at http://webbook.nist.gov/chemistry.

Compound

S° ( J mol–1 K–1)

Compound

S° ( J mol–1 K–1)

H2(g)

130.6

CO2(g)

213.6

O2(g)

205.0

C4H10(g)

310.03

CH4(g)

186.2

69.91

H2O(,) H2O(g)

188.7

C2H4(g)

219.5

NH3(g)

192.3

C3H3N(,)

178.91

convenient method comes from the third law of thermodynamics, which says that the entropy of a perfect crystal of any pure substance approaches zero as the temperature approaches absolute zero. An additional implication of the third law is that it is impossible to attain a temperature of absolute zero, although scientists have come very close to that value. Because all substances can be cooled to temperatures near zero (at least in principle), it is possible to evaluate the entropy of one mole of any given chemical substance under standard conditions, by determining the change in entropy from 0 K to 298 K at a pressure of 1 atm. This process yields the standard molar entropy, S°. Table 10.1 provides S° values for some substances. A more complete list is given in Appendix E. Because entropy is a state function and because the third law allows us to obtain a value for the standard molar entropy of any substance, we can derive a useful equation for the entropy change in a reaction. Figure 10.5 shows how the entropy change in a reaction may be determined by a method that is reminiscent of the way we used heats of formation and Hess’s law in Chapter 9. DS° = ∑ ni S°( product)i 2 ∑ nj S°(reactant)j i

j

(10.3)

Here, we designate the stoichiometric coefficient by the Greek letter n, and the subscripts i and j refer to individual product and reactant species. In practice, this equation is used much like Hess’s law, as shown in the following example. One subtle difference between this and Equation 9.12 is that here we use absolute entropy values (S°) in contrast to the enthalpy changes for formation reactions (DHfº).

The entropy change for a reaction may not be easy to determine directly. Products Reactants Entropy

Figure 10.5 ❚ Entropy is a state function, so the value of DS must be independent of the path taken from reactants to products. When working with tabulated standard molar entropies, we implicitly choose a path in which reactants are converted to pure perfect crystals of elements at 0 K, followed by the reaction of those elements to form the desired products. Such a path is obviously not feasible. Nevertheless, it allows us to obtain accurate values of entropy changes.

This portion of the path corresponds to the opposite of the tabulated S° values of the reactants.

This portion of the path corresponds to the tabulated S° values of the products. Elements in perfect crystals at 0 K

Because entropy is a state function the value of ΔS for the two paths is the same.

328



10 .1

Polymerization reactions are complicated somewhat because they involve very large numbers of molecules. But we can demonstrate the general features of the thermodynamics of polymerization by considering a much smaller model system. Instead of considering the formation of polyethylene, for example, we can begin with the following reaction in which two ethylene molecules combine with hydrogen to form butane: 2 C2H4(g) 1 H2(g) : C4H10(g) Use data from Table 10.1 to calculate DSº for this reaction.

Strategy Any time we are asked to calculate the standard entropy change for a reaction, our first thought should be to look up values for standard molar entropy and use them in Equation 10.3. The two main things we need to be careful about are (1) to watch the state of the substances (in this case all are gases) and (2) to make sure we don’t forget to include the stoichiometric coefficients in our calculations. Unlike heats of formation, the standard molar entropy of an element in its standard state is not zero, so we need to be sure to include everything appearing in the equation.

Solution DS° 5 S°[C4H10(g)] 2 2S°[C2H4(g)] 2 S°[H2(g)] 5 (310.03 J K21) 2 2(219.5 J K21) 2 (130.6 J K21) 5 2259.6 J K21

Analyze Your Answer You probably don’t have an intuition for the size of an entropy change, but at least we can think about the sign. Does it make sense that this reaction has a negative change of entropy? The answer is yes, because we have decreased the amount of gas present as a result of this reaction. Three moles of gaseous reactants are consumed, and only one mole of gaseous product is formed.

Discussion Can this reaction be spontaneous? Yes, as it turns out. The explanation lies in the change in entropy of the surroundings. Because it forms a number of strong C!H and C!C bonds, this reaction is exothermic. The release of heat will increase the entropy of the surroundings. This factor must be large enough to compensate for the decrease in entropy of the system itself.

Check Your Understanding Acrylonitrile (C3H3N) is an important monomer for the production of many acrylic fibers. It can be synthesized from propene and ammonia according to the following reaction: 2 C3H6(g) 1 2 NH3(g) 1 3 O2(g) : 2 C3H3N(,) 1 6 H2O(g) Given that DS° for this reaction is 243.22 J/mol, use data from Table 10.1 to calculate the standard molar entropy (S°) of C3H6(g).

Example Problem 10.1 shows one of the dangers of attempting to use DS to predict spontaneity. The need to consider both the system and surroundings creates additional work, and it also allows drawing an incorrect conclusion if we forget to include the surroundings. As you might guess, having to account for the surroundings is rarely convenient. After all, by definition, the system is what we are really interested in! Ideally we would be able to use a state function that predicts spontaneity with a single calculation for the system alone. Fortunately, such a state function exists, and so we now proceed to introduce the concept of free energy.

10.5


329

10.6 An alternative function called the Helmholtz free energy is useful for constant volume conditions. This is less common in engineering applications, so we will not consider it.

Gibbs Free Energy

In many ways, thermodynamics is one of the more powerful examples of the application of mathematics to science. If you continue in engineering, there is a good chance that you will eventually take a full course in thermodynamics in which you will develop a much more rigorous mathematical perspective. For now, we simply will introduce the fruits of the mathematical efforts of one of the scientists who invented thermodynamics, J. Willard Gibbs. Gibbs, motivated by an interest in predicting spontaneous processes, defined a new function that was ultimately called the Gibbs free energy, G: G 5 H 2 TS He realized that changes in this function could predict whether or not a process is spontaneous under conditions of constant pressure and temperature. This constraint is not that demanding because many laboratory processes occur under these conditions (or approximately so). Now, if we focus on the change in the Gibbs free energy for a process at constant temperature, we have the following result: DG 5 DH 2 TDS

(10.4)

From a practical viewpoint, this could be the most important equation in this chapter.

Free Energy and Spontaneous Change How is this change in free energy connected to the spontaneity of a process? We have already established that the total entropy of the system plus its surroundings must increase for a spontaneous process. We can write that total entropy change as follows: DSu 5 DSsys 1 DSsurr We usually refer to the entropy change in the system simply as DS, so we’ll drop the “sys” subscript. DSu 5 DS 1 DSsurr The entropy change of the surroundings is due to the flow of heat, and because T is constant, it can be shown that DSsurr is just qsurr/T. And qsurr is 2qsys. For a process at constant pressure, qsys is DH. Combining these ideas, we have now established that DH DSsurr 5 2 —— T We can insert that into our equation above for DSu. DH DSu 5 DS 2 —— T If we multiply both sides of this by T, we should begin to see a connection to DG. TDSu 5 TDS 2 DH Comparing this with DG 5 DH 2 TDS we can see that DG 5 2TDSu In all of these equations, T is the absolute temperature, so its value is always positive. That means that the last equation ensures that the sign of DG will always be opposite that of DSu. So if DSu is positive—as it must be for any spontaneous process—then DG must be negative.

330


Table

❚

10.2

The four possible combinations for the signs of DH and DS

Implications for Spontaneity

Sign of DH

Sign of DS

–

+

Spontaneous at all temperatures

+

–

Never spontaneous

–

–

Spontaneous only at low temperatures

+

+

Spontaneous only at high temperatures

All of this shows why chemists find DG such a useful thermodynamic quantity. It is a state function of the system, so it can be calculated fairly easily. The sign of DG is sufficient to tell us whether or not a process will be spontaneous. We can also use Equation 10.4 to help formalize our understanding of the roles of DH and DS in determining the spontaneity of a given reaction. We have already argued that exothermic reactions, with DH < 0, seem to be preferred over endothermic ones and also that reactions where DS is positive seem to be preferred. Equation 10.4 shows us that if DH is negative and DS is positive, then DG will always be negative. But not all spontaneous processes fit this specific pattern. Table 10.2 shows the four possible combinations of signs for DH and DS. If a process is exothermic, but entropy decreases, we see that the sign on DG depends on temperature. As T increases, the relative importance of the 2TDS term also increases, so such processes will be spontaneous only at lower temperatures where the DH term is dominant. Processes that occur spontaneously only at lower temperatures are sometimes said to be enthalpy driven because the enthalpy term is responsible for the negative value of DG. For endothermic processes, if the entropy of the system decreases, the sign on DG will always be positive, and the process is never spontaneous. An endothermic process that increases the entropy, however, may be spontaneous at high temperatures, where the 2TDS term becomes larger than the DH term of Equation 10.4. These processes are said to be entropy driven. The reasoning associated with Table 10.2 can be used to understand the nature of phase changes, as noted in Example Problem 10.2.


The word “driven” is used here to imply that either the entropy or enthalpy term is dominant in determining the sign of DG. There is no force due to enthalpy or entropy responsible for driving spontaneous changes.

10 . 2

Use the signs of DH and DS to explain why ice spontaneously melts at room temperature but not outside on a freezing winter day.

Strategy This problem calls for the same type of reasoning used to construct Table 10.2. We must determine whether the process is endothermic or exothermic and whether it increases or decreases the entropy of the system. Then, by considering the signs of DH and DS in conjunction with Equation 10.4, we can attempt to explain the behavior described.

Solution Melting is an endothermic process (DH . 0) because we must heat the system to effect the change. It is also a process that increases entropy because molecules formerly held in place in a solid have greater freedom of motion in a liquid and are therefore less ordered. Thus, both DH and DS have positive values, and DG will be negative at high temperatures, where the TDS term is larger. So, melting tends to occur at higher temperatures. At least for water, we know that room temperature is

10.6 Gibbs Free Energy

331

sufficient to make melting spontaneous. At low temperatures, the DH term is more important, so on a freezing day, the sign of DG is positive. Melting is not spontaneous and therefore is not observed. At the freezing point, DG is equal to zero, and ice and water can coexist in any proportions.

Discussion The result that ice melts if it is warm but not if it is cold is intuitively obvious. But this problem emphasizes the importance of gaining a qualitative understanding of entropy changes.

Check Your Understanding At what temperatures would you expect a gas to condense? Explain why the free energy change is negative for this process at these temperatures.

If the actual values of DH and DS are known, then the same type of argument used in Example Problem 10.2 can also be extended to obtain quantitative information.

E XAM PL E PRO BLEM

10 .3

The heat of fusion of crystalline polyethylene is approximately 7.7 kJ/mol, and the corresponding entropy change for melting is 19 J/mol K. Use these data to estimate the melting point of polyethylene.

Strategy Because DH and DS are both positive, we know that DG must be positive at low temperatures and negative at higher temperatures. The melting point marks the dividing line between low and high temperatures. So at the melting point itself, DG must be equal to zero. We can therefore start with Equation 10.4 and set DG 5 0. That leaves T as the only unknown, so we can solve for the desired melting point.

Solution We begin with Equation 10.4 and set DG 5 0. (Adding a subscript ‘m’ to the temperature will help remind us that this equation is only valid at the melting temperature.) DG 5 DH 2 TmDS 5 0 Rearranging this gives us DH 5 Tm DS Because DH and DS are known, solving for Tm is simple. (We do need to take care in handling the units, though, because DH is in kJ and DS is in J.) 7700 J mol21 DH Tm 5 —— 5 ——————— 5 410 K DS 19 J mol21 K21

Analyze Your Answer The result seems reasonable: 410 K is 140°C, or 280°F, which is a plausible melting point for a plastic. Note that if we had been sloppy with units and forgotten to convert DH from kJ/mol into J/mol we would have gotten an answer of 0.4 K. Such an extremely low temperature should then have tipped us off to the mistake.

Discussion You may recall from our earlier discussions of polymers that the molecules in a given polymer usually do not all have the same chain length. Thus there is a bit of difficulty in specifying quantities per mole. The values used here were actually measured per mole of monomer unit.

332


Check Your Understanding Poly(tetrafluoroethylene) melts at approximately

327°C. If the heat of fusion is DHfus 5 5.28 kJ mol21, what is the molar entropy change of fusion (DSfus)?

Free Energy and Work By now, you are probably wondering about the name. Why is it called free energy? The change in Gibbs free energy can be shown to be equal to the maximum useful work that can be done by the system, DG 5 2 wmax

(10.5)

We must include a minus sign in this expression to be consistent with the convention that w is the work done on the system. This relationship suggests that DG tells how much energy is “free” or available to do something useful. For an engineer interested in making practical use of a chemical reaction, the implications should be clear. Keep in mind that work is not a state function. So the maximum work will be realized only if we carry out the process by a specific path. In this case, the requirement is that the change is carried out along a reversible path. This means that the system is near equilibrium, and a small incremental change in a variable will bring the system back to its initial state. Maximum work is possible only for reversible processes. Systems that are far from equilibrium usually undergo irreversible changes. In an irreversible change, a small incremental change of any variable does not restore the initial state. The amount of work available in an irreversible change is always less than the maximum work. So, although the free energy change can be used to establish an upper bound to the amount of work obtained from a given process, the actual work produced in any real application may be considerably less. Reactant mixtures such as those for combustion reactions are generally very far from equilibrium. Systems that are far from equilibrium often change rapidly, and rapid changes tend to be irreversible. To use these free energy concepts quantitatively, of course, we still need a simple method for obtaining accurate values of free energy changes.

10.7

Calculations for irreversible changes are challenging, and we will not attempt any such exercises.

Free Energy and Chemical Reactions

Free energy is a state function, and so the value of the free energy of a system depends on specific variables such as concentration or pressure. As usual, to provide consistent comparison, we define a standard state as 1 atmosphere of pressure and concentrations of solutions of 1 M. The free energy change under these conditions is the standard Gibbs free energy change, DG°. Although it is feasible to calculate this value from DH° and S° at a given temperature using Equation 10.4, the most convenient means for calculating the change in free energy for many reactions is to use a formulation analogous to Hess’s law for enthalpy changes. DG° = ∑ ni Gf°( product)i 2 ∑ nj Gf°(reactant)j i

j

Hess’s law relating enthalpy changes to terms of heats of formation was given in Equation 9.12, page 304.

(10.6)

Again in Equation 10.6, we use the concept of the formation reaction from Chapter 9. Tabulated values of free energies of formation for a few substances are provided in Table 10.3, and a more extensive list appears in Appendix E. Note that DGf° is zero for elements in their standard states, for the same reason that their heats of formation are zero. Because the formation reaction uses elements in their standard states to define the reactants, a formation reaction of an element would have the same chemical species as both reactant and product, and this process would clearly have no change in any thermodynamic state function. We can use Equation 10.6 to calculate standard free energy changes, as demonstrated in Example Problem 10.4.

10.7 Free Energy and Chemical Reactions

333

Table

❚

10.3

Values of the free energy change of formation, DGf°, for selected compounds. A much larger list appears in Appendix E.

Compound

DGf° (kJ mol–1)

Compound

DGf° (kJ mol–1)

H2(g)

0

CO2(g)

–394.4

O2(g)

0

C4H10(g)

–15.71

H2O(,)

–237.2

CH4(g)

–50.75

H2O(g)

–228.6

C2H4(g)

68.12

NH3(g)

–16.5

C3H6(g)

62.75

E XAM PL E PRO BLEM

10 .4

In Example Problem 10.1, we considered an addition reaction involving two ethylene molecules and found that the entropy change was negative. We suggested at the time that this reaction would still be spontaneous because it is strongly exothermic. Confirm this by calculating the standard free energy change for the same reaction using values from Table 10.3. 2 C2H4(g) 1 H2(g) : C4H10(g)

Strategy Any problem that requests a calculation of a state function using values from a table generally means we’ll be using a formulation like Equation 10.6. Free energy is the third such state function we have encountered that can be treated in this way. If we’re careful to note the states of the molecules present and watch the stoichiometric coefficients, this type of problem can be solved readily.

Solution DG° 5 DGf°[C4H10(g)] 2 2DGf°[C2H4(g)] 2 DGf°[H2(g)] 5 (215.71 kJ) 2 2(68.12 kJ) 2 (0) 5 2151.95 kJ

Analyze Your Answer This value is negative, indicating that this reaction would be spontaneous under standard conditions, as we had said. Many reactions have free energy changes on the order of 102 kJ/mol, so our result seems plausible.

Discussion It’s worth pointing out again, though, that this does not mean that the reaction will occur readily if we mix the reactants. Simply because a reaction has a negative value for DG° does not mean we will observe a spontaneous conversion. In some cases, the rate of the reaction is so slow that, despite the thermodynamic indication of spontaneity, the reaction is not observed. Thermodynamics does not tell us how rapidly a spontaneous process will take place.

Check Your Understanding For the reaction shown below, DG° 5 21092.3 kJ. Find DGf° for liquid acrylonitrile (C3H3N). 2 C3H6(g) 1 2 NH3(g) 1 3 O2(g) : 2 C3H3N(,) 1 6 H2O(g)

334


Implications of DG° for a Reaction Now that we have a method for calculating the standard free energy change, what does this value tell us? One answer is that the free energy change tells us the maximum useful work that can be obtained from a reaction. But the free energy change also has very important implications for the many chemical reactions that can be made to run in either direction. Earlier in this chapter, we considered the thermal depolymerization (or thermolysis) of PMMA as a possible alternative to mechanical recycling. Now, we can use free energy changes to explore that idea quantitatively. The polymerization of methyl methacrylate has DH° 5 256 kJ and DS° 5 2117 J/K. We can use these values in Equation 10.4 to find DG° 5 221 kJ. (We use the standard temperature of 298 K, which gives us the standard free energy change.) The negative value tells us that the formation of the polymer is spontaneous at 298 K. This, of course, means that the reverse reaction in which PMMA is broken down into monomers cannot be spontaneous at this temperature. Because the depolymerization amounts to running the polymerization reaction backward, we can conclude that it has DG° 5 121 kJ. Polymerization: methyl methacrylate monomer : PMMA polymer DG° 5 221 kJ Depolymerization: PMMA polymer : methyl methacrylate monomer DG° 5 121 kJ The fact that the numerical values of these free energy changes are relatively small hints at the fact that a relatively small shift in temperature might flip the signs, making depolymerization the thermodynamically preferred process. We can quantify this idea, too. Using the same sort of calculation as we did in Example Problem 10.3, we can find the temperature at which DG° will equal zero. 256 kJ DH T 5 —— 5 ——————— 5 480 K DS 20.117 kJ K21 Above this temperature, entropy wins out over enthalpy, and the reaction will be driven backward, toward the monomer. The fact that this temperature threshold is reasonably low is one reason why depolymerization is a viable recycling strategy for PMMA. Most polymerization reactions have DS° values around 2100 J/K, but many are more strongly exothermic than that for PMMA. Looking at the equation above, we can see that this leads to a higher temperature requirement for depolymerization. So thermolysis is less feasible for two reasons. First, the need for higher temperatures implies higher cost. Second, the higher temperatures increase the likelihood that additional chemical reactions, such as breakdown of monomers into other compounds, would compete with thermolysis.

INSIGHT INTO

10.8

The Economics of Recycling

According to the Container Recycling Institute, about 45% of all aluminum beverage cans sold in the United States in 2006 were recycled. But in the same year only about 24% of PET bottles were recycled. It seems fair to assume that individual consumers are no more anxious to recycle aluminum cans than plastic bottles. So there must be a real cause for the difference between those two numbers, and that cause lies in economics. When it comes right down to it, the business of recycling is all about trying to sell your trash. And as you might guess, selling trash can be a difficult business. In this section we’ll take a look at some of the factors that make aluminum recycling so much more attractive than plastic recycling.

10.8

Recycling rates vary significantly from place to place, and are significantly higher in states with deposit laws. The numbers cited here are national averages.


335

Aluminum foil made entirely from recycled aluminum has also been introduced recently.

336

We’ll begin by looking at the recycling of aluminum, which now accounts for nearly all beverage cans sold in the United States. The chemistry of aluminum makes it an excellent target for recycling. Recall that in Chapter 1 we pointed out how difficult it is to obtain pure aluminum from ores like bauxite. Aluminum reacts readily with oxygen, forming strong chemical bonds that are not easily disrupted. Extracting aluminum from its ores requires very high temperatures and therefore consumes a lot of energy. When aluminum cans are recycled, however, the paint and other coatings used can be removed relatively easily, allowing the underlying aluminum to be melted down and reprocessed into a new can. Current industry estimates are that making four new cans from recycled aluminum uses the same amount of energy as producing one can from raw aluminum ore. This provides strong economic incentives throughout the entire process. First, the cost savings mean that the beverage industry has a strong motivation to encourage recycling and to seek recycled aluminum to produce new cans. This, in turn, means that communities and private companies in the recycling business are assured of finding a good market for any aluminum cans they can collect. This is an important concern because there will always be costs for collecting recyclables. Other positive factors also exist in the recycling equation for aluminum. The fact that virtually all beverage cans in use are made of aluminum means there is no need to separate collected cans into different types. And because aluminum cans are so easily crushed, the collected material can be compressed so that storage and transportation are easier and less expensive. All in all, aluminum is an ideal candidate for recycling. Now let’s contrast that with plastic recycling. We can begin by thinking about the production costs of virgin polymers because these costs will set the standard against which the price of recycled plastics will be measured. Feedstocks for virtually all commercial polymerization reactions have their roots in petroleum. So, the cost of raw materials for synthesizing most plastics is linked to the price of oil. Oil is a complex mixture of compounds, and our new understanding of entropy should make clear that such a mixture is unlikely to separate spontaneously into its various components. Various separation and purification methods are used to obtain the needed monomer molecules. Most of these schemes involve cracking and distillation, which require heating the crude oil until various components boil off and can be reclaimed from the vapor phase. This need for heating means that there is also an energy cost for producing the feedstock for polymerization. Once we have a supply of the appropriate monomer (and any other reagents that may be needed), we will have to pay to transport them to the plant where the polymer will be produced. That cost can be minimized, of course, if the plant where the polymer will be made is located close to a refinery where the oil is processed. Many polymerization reactions are spontaneous under ordinary conditions but as we’ve learned, that doesn’t necessarily mean that those reactions are fast. If we are in the business of making plastics, we will probably not be satisfied if our polymer production requires days or even years. Most reactions run faster when heated, though, as we’ll see when we examine chemical kinetics in Chapter 11. So, we will probably want to carry out our polymerization at higher temperatures. This will add a further energy cost to the bill for producing our polymer. So how will this compare to the cost of recycling plastics? Just as for aluminum cans, there will be some costs for collecting the bottles to be recycled. But the fact that plastic bottles are made from a wide array of polymers offers some added complications here. The various types of polymers usually must be separated before they can be processed. Consumers can be encouraged to do some of this separation based on the recycling codes commonly found on bottles (see Table 10.4). But even different colors of the same type of plastic may be incompatible, and inevitably there will be some mixing of bottle types in the collection bins. So the recycler must expect to have to sort the materials before they can be sold. This is most often done by hand, although sometimes flotation and separation based on density may be possible. No matter the method, though, this separation adds cost to the overall recycling effort. Once bottles of a particular type have been separated, they are ready to be processed.


Table

❚

10.4

Symbols, structures, sources, and uses for various recycled plastics

Symbol

1

Polymer

Repeat Unit

Polyethylene terephthalate

*

[ C

PETE

O

2

O

O

High-density polyethylene

* n

*

C O

Sources

Recycled Products

Soda bottles, peanut butter jars, vegetable oil bottles

Fiber, tote bags, clothing, film and sheet, food and beverage containers, carpet, fleece wear

Milk and water jugs, juice and bleach bottles

Bottles for laundry detergent, shampoo, and motor oil; pipe, buckets, crates, flower pots, garden edging, film and sheet, recycling bins, benches, dog houses, plastic lumber

Detergent / cleanser bottles, pipes

Packaging, loose-leaf binders, decking, paneling, gutters, mud flaps, film and sheet, floor tiles and mats, resilient flooring, electrical boxes, cables, traffic cones, garden hose

Six-pack rings, bread bags, sandwich bags

Shipping envelopes, garbage can liners, floor tile, furniture, film and sheet, compost bins, paneling, trash cans, landscape timber, lumber

Margarine tubs, straws, screw-on lids

Automobile battery cases, signal lights, battery cables, brooms, brushes, ice scrapers, oil funnels, bicycle racks, rakes, bins, pallets, sheeting, trays

Styrofoam, packing peanuts, egg cartons, foam cups

Thermometers, light switch plates, thermal insulation, egg cartons, vents, desk trays, rulers, license plate frames, foam packing, foam plates, cups, utensils

Squeezable ketchup and syrup bottles

Bottles, plastic lumber

[n*

HDPE

3

Polyvinyl chloride

PVC

4

Low-density polyethylene

Cl * *

n

*

n

*

LDPE

Polypropylene

CH3

5

* *

PP

n

Polystyrene

6 PS

* *

7

Miscellaneous and multilayer plastics

N/A

n

Other

This would typically include crushing, washing, and melting the plastic before forming it into pellets to be resold. Again, each of those steps adds some cost. Although consumers may toss a bottle into a recycling bin out of a sense of doing a good deed, a company that might be interested in buying the recycled plastic pellets is more likely to have its eye on the bottom line. For the resulting recycled pellets to be attractive, they will have to be priced competitively with virgin polymer. And here things are not nearly as favorable as for aluminum. The production of most plastics from raw feedstocks is generally less expensive than the processing of aluminum from

10.8


337

Recall that ordinary samples contain a range of polymer chain lengths. The degradation associated with recycling shifts this distribution toward shorter chains.

its ore. So, in economic terms, the standard to which recycled plastic will be held is more demanding. There is also an important difference in the ability of recycling to produce high quality materials. Once the aluminum has been reclaimed from a recycled can and melted down, it is indistinguishable from new aluminum freshly extracted from ore. But for plastics, the recycling process leads to degradation of the polymer. This is really just an effect of the second law: the entropy of a polymer molecule will generally increase if the long chain is broken into shorter pieces. Chain lengths in recycled plastics invariably are shorter than in virgin materials, and this can raise concerns as to whether the recycled material offers sufficient strength to satisfy design requirements. On balance, then, many businesses find it less expensive to use virgin plastics rather than recycled polymers. Are there circumstances that would be likely to shift this balance? That would require either that recycling become less expensive or that the production of virgin polymers become more expensive, or both. The cost of virgin materials has its roots in the price of the petroleum from which the needed feedstocks are obtained. So a significant rise in oil prices would tend to drive up the cost of nearly all plastics, too. But because many of the costs of recycling are energy related, they are also tied in one way or another to oil prices. So it is not immediately clear that plastic recycling can be made more economically attractive in the near term. If recycled plastics cannot compete effectively with virgin materials in applications such as drink bottles, then might there be other outlets in which the material can be put to use? One increasingly popular idea is to use recycled plastic in applications typically calling for more traditional (i.e., nonplastic) materials. The rapidly growing market for “lumber” made from recycled plastic is the most prominent example. Recycled plastic is now fairly widely used in building decks, picnic tables, playground equipment, and other items that we generally picture as made of wood. Although this may not appeal to traditionalists, the plastic alternatives offer better weather resistance and durability. Thus they can be a cost-saving choice, especially if one considers long-term savings on maintenance. Other areas in which recycled plastic now competes with nonplastic materials include filling for winter coats and sleeping bags.

FOCU S O N

PRO BLEM S O LVI NG

Question Suppose that you need to know the melting point of an oil for the design of a microwaveable food package. You have no sample of the oil, so you can’t measure the melting point. But you can find tabulated thermodynamic data for the oil in both solid and liquid phases. What specific values would you need to look up in the table, and how would you use them to determine the melting point of the oil?

Strategy This question asks us to think about how a phase change is described thermodynamically, and how we can use tabulated data along with the equations to get the melting point. It will help if we think of melting the oil as a simple reaction: Oil(s) : Oil(,)

Answer From Example Problem 10.3, we know that the melting point is given by the ratio of DH/DS. If we have data for both the solid and the liquid forms of the oil, we can determine the enthalpy of fusion (DHfus) by subtracting the heat of formation of the solid from that of the liquid. Similarly, we can get the entropy change for fusion (DSfus) by subtracting the absolute entropy of the solid from that of the liquid. Taking the ratio of these two numbers will tell us the melting temperature.

338


S U M M ARY Thermodynamics can provide more information than just the amount of energy released or absorbed in a chemical or physical process, as discussed in Chapter 9. By defining additional state functions, we can also determine the spontaneous direction of change of any system—a very powerful tool for understanding chemistry and an important factor in many engineering designs. To predict spontaneity, we introduced two new concepts and state functions: entropy and free energy. We can define entropy in either of two ways: as the ratio of the heat flow to the temperature or as a measure of the number of ways that a system can have the same energy. This latter definition, for practical purposes, is a measure of the extent of the randomness of the system

at the atomic and molecular level. A more random system will have more ways to distribute its energy, so entropy increases with the extent of randomness. The second law of thermodynamics tells us that spontaneous changes always increase the entropy of the universe. But calculations that must consider the universe are rarely practical. So we define another state function, called Gibbs free energy, so that its change predicts spontaneity. The change in Gibbs free energy, given by DG 5 DH 2 TDS, is always negative for a spontaneous process. Both entropy changes and free energy changes can be calculated for many chemical and physical processes by using tabulated thermodynamic data.

K EY TE RM S enthalpy driven (10.6)

microstate (10.3)

standard molar entropy (10.5)

entropy (10.3)

reversible (10.6)

statistical mechanics (10.3)

entropy driven (10.6)

second law of thermodynamics (10.4)

third law of thermodynamics (10.5)

Gibbs free energy (10.6)

spontaneous process (10.2)

irreversible (10.6)

standard Gibbs free energy change (10.7)



INSIGHT INTO Recycling of Plastics 10.1 “Reduce, reuse, recycle” is a common slogan among environmentalists, and the order of the three words indicates their perceived relative benefits. Why is recycling the least desirable of these three approaches to waste reduction? 10.2 Is the recycling of most plastics primarily a chemical or a physical process? Explain and defend your choice. 10.3 List some consumer products made from recycled PET. 10.4 Why is recycled PET rarely used to make new soft drink bottles? 10.5 Use the web to research a company that specializes in the recycling of plastics. Does the material on their website emphasize environmental, scientific, or economic concerns? Write a brief essay on the company’s positions, explaining how they fit with the ideas expressed in this chapter. 10.6 Use the web to learn how many pounds of plastics are recycled in your area each year. How has this value changed during the past decade?

■


Spontaneity 10.7

On the basis of your experience, predict which of the following reactions are spontaneous.

■

(a) CO2(s) : CO2(g) at 25°C (b) NaCl(s) : NaCl(,) at 25°C (c) 2 NaCl(s) : 2 Na(s) 1 Cl2(g) (d) CO2(g) : C(s) 1 O2(g) 10.8 In the thermodynamic definition of a spontaneous process, why is it important that the phrase “continuous intervention” be used rather than just “intervention?” 10.9 If the combustion of butane is spontaneous, how can you carry a butane lighter safely in your pocket or purse? 10.10 Identify each of the processes listed as spontaneous or nonspontaneous. For each nonspontaneous process, describe the corresponding spontaneous process in the opposite direction. (a) A group of cheerleaders builds a human pyramid. (b) Table salt dissolves in water. (c) A cup of cold coffee in a room becomes steaming hot.


339

I ever did fall off—which there’s no chance of—but if I did . . . the King has promised me—with his very own mouth—(that) they’d pick me up again in a minute, they would!”

(d) Water molecules in the air are converted to hydrogen and oxygen gases. (e) A person peels an orange, and you smell it from across the room.

Write a paragraph in the voice of seven-and-a-halfyear-old Alice, explaining to Humpty in the context of this section (a) the probability that Humpty will fall off the wall and (b) the probability that the King’s horses and men will be able to put him back together again.

10.11 Identify each of the processes listed as spontaneous or nonspontaneous. For each nonspontaneous process, describe the corresponding spontaneous process in the opposite direction. (a) Oxygen molecules dissociate to form oxygen atoms. (b) A tray of water is placed in the sun on a warm day and freezes.

10.18

(c) A solution of salt water forms a layer of acid on top of a layer of base. (d) Silver nitrate is added to a solution of sodium chloride and a precipitate forms. (e) Sulfuric acid sitting in a beaker turns into water by giving off gaseous SO3.

The vessel on the left contains a mixture of oxygen and nitrogen at atmospheric pressure. The vessel on the right is evacuated.

10.12 Athletic trainers use instant ice packs that can be cooled quickly on demand. Squeezing the pack breaks an inner container, allowing two components to mix and react. This reaction makes the pack become cold. Describe the heat flow for this spontaneous process.

(a) Describe what will happen when the stopcock is opened. (b) If you could see the individual molecules, what would you observe after a period of time has passed?

10.13 Are any of the following exothermic processes not spontaneous under any circumstances?

(c) Explain your answers to (a) and (b) in terms of probabilities.

(a) Snow forms from liquid water.

(d) What is the probability that at any one moment all the oxygen molecules will be in one vessel and all the nitrogen molecules will be in the other? Explain.

(b) Liquid water condenses from water vapor. (c) Fossil fuels burn to form carbon dioxide and water. (d) Monomers react to form a polymer. 10.14 Enthalpy changes often help predict whether or not a process will be spontaneous. What type of reaction is more likely to be spontaneous: an exothermic or an endothermic one? Provide two examples that support your assertion and one counterexample. 10.15 When a fossil fuel burns, is that fossil fuel the system? Explain your answer. 10.16 Murphy’s law is a whimsical rule that says that anything that can go wrong will go wrong. But in an article in the Journal of Chemical Education, Frank Lambert writes, “Murphy’s law is a fraud.” He also writes, “The second law of thermodynamics is time’s arrow, but chemical kinetics is its clock.” Read Lambert’s article ( J. Chem. Ed., 74(8), 1997, p. 947), and write an essay explaining, in the context of the latter quotation, why Lambert claims that Murphy’s law is a fraud. (For more of Professor Lambert’s unique insights into thermodynamics, see his website at http://www.secondlaw.com/) 10.17

Humpty Dumpty sat on a wall, Humpty Dumpty had a great fall. All the King’s horses and all the King’s men Couldn’t put Humpty together again. In Lewis Carroll’s Through the Looking Glass, Alice encounters Humpty Dumpty, a human-sized egg sitting on a wall. Alice, who is familiar with the nursery rhyme, asks anxiously, “Don’t you think you’d be safer on the ground? That wall is so narrow.” Humpty, an egg with an attitude, growls, “Of course I don’t think so. Why if

340

Entropy 10.19 What observation about the Carnot engine led Carnot to propose the existence of a new state function? 10.20 Some games include dice with more than six sides. If you roll two eight-sided dice, with faces numbered one through eight, what is the probability of rolling two eights? What is the most probable roll? 10.21 How does probability relate to spontaneity? 10.22 Define the concept of a microstate. How is this concept related to the order or disorder of a system? 10.23

■ For each pair of items, tell which has the higher entropy and explain why.

(a) Item 1, a sample of solid CO2 at 278°C, or item 2, CO2 vapor at 0°C (b) Item 1, solid sugar, or item 2, the same sugar dissolved in a cup of tea (c) Item 1, a 100-mL sample of pure water and a 100-mL sample of pure alcohol, or item 2, the same samples of water and alcohol after they have been poured together and stirred 10.24 When ice melts, its volume decreases. Despite this fact, the entropy of the system increases. Explain (a) why the entropy increases and (b) why under most circumstances, a decrease in volume results in an entropy decrease. 10.25 If a sample of air were separated into nitrogen and oxygen molecules (ignoring other gases present), what would be the sign of DS for this process? Explain your answer.


■


10.26

10.27

For each process, tell whether the entropy change of the system is positive or negative. (a) A glassblower heats glass (the system) to its softening temperature. (b) A teaspoon of sugar dissolves in a cup of coffee (the system consists of both sugar and coffee). (c) Calcium carbonate precipitates out of water in a cave to form stalactites and stalagmites. (Consider only the calcium carbonate to be the system.)

■

Without doing a calculation, predict whether the entropy change will be positive or negative when each of the following reactions occurs in the direction it is written.

■

(a) CH3OH(,) 1 3/2 O2(g) : CO2(g) 1 2 H2O(g) (b) Br2(,) 1 H2(g) : 2 HBr(g) (c) Na(s) 1 1/2 F2(g) : NaF(s) (d) CO2(g) 1 2 H2(g) : CH3OH(,) (e) 2 NH3(g) : N2(g) 1 3 H2(g) 10.28

For the following chemical reactions, predict the sign of DS for the system. (Note that this should not require any detailed calculations.)

■

(a) Fe(s) 1 2 HCl(g) : FeCl2(s) 1 H2(g) (b) 3 NO2(g) 1 H2O(,) : 2 HNO3(,) 1 NO(g) (c) 2 K(s) 1 Cl2(g) : 2 KCl(s) (d) Cl2(g) 1 2 NO(g) : 2 NOCl(g) (e) SiCl4(g) : Si(s) 1 2 Cl2(g) 10.29 In many ways, a leaf is an example of exquisite order. So how can it form spontaneously in nature? What natural process shows that the order found in a leaf is only temporary? The next four questions relate to the following paragraph (Frank L. Lambert, Journal of Chemical Education, 76(10), 1999, 1385). “The movement of macro objects from one location to another by an external agent involves no change in the objects’ physical (thermodynamic) entropy. The agent of movement undergoes a thermodynamic entropy increase in the process.” 10.30 A student opens a stack of new playing cards and shuffles them. In light of the paragraph above, have the cards increased in entropy? Explain your answer in terms of thermodynamics. Explain why the agent (the shuffler) undergoes an increase in entropy. 10.31 An explosion brings down an old building, leaving behind a pile of rubble. Does this cause a thermodynamic entropy increase? If so, where? Write a paragraph explaining your reasoning. 10.32 Write two examples of your own that illustrate the concept in the paragraph above. 10.33 According to Lambert, leaves lying in the yard and playing cards that are in disarray on a table have not undergone an increase in their thermodynamic entropy. Suggest another reason why leaves and playing cards may not be a good analogy for the entropy of a system containing, for example, only H2O molecules or only O2 molecules. 10.34 A researcher heats a sample of water in a closed vessel until it boils. (a) Does the entropy of the water increase? (b) Has the randomness of the molecules increased? (In other words, are there more physical positions that the molecules can occupy?)

■


(c) What else has increased that affects the entropy of the system? The researcher now heats the water vapor from 400 K to 500 K, keeping the volume constant. (d) Does the entropy of the system increase? (e) Has the randomness of the molecules increased? (In other words, are there more physical positions that the molecules can occupy?) (f) Why has an increase in temperature of the gas at constant volume caused an increase in entropy? (To delve more deeply into the concept of entropy, read John P. Lowe’s article “Entropy: Conceptual Disorder” in the Journal of Chemical Education, 65(5), 1988, 403– 406.) The Second Law of Thermodynamics 10.35 What happens to the entropy of the universe during a spontaneous process? 10.36 Why do we need to consider the surroundings of a system when applying the second law of thermodynamics? 10.37 One statement of the second law of thermodynamics is that heat cannot be turned completely into work. Another is that the entropy of the universe always increases. How are these two statements related? 10.38 According to the second law of thermodynamics, how does the sign of DS u relate to the concept that some energy is wasted or lost to the surroundings when we attempt to convert heat into work? 10.39 How does the second law of thermodynamics explain a spontaneous change in a system that becomes more ordered when that process is exothermic? 10.40 Some say that the job of an engineer is to fight nature and the tendencies of entropy. (a) Does this statement seem accurate in any way? (b) How can any engineering design create order without violating the second law of thermodynamics? 10.41 When a reaction is exothermic, how does that influence DS of the system? Of the surroundings? 10.42 Which reaction occurs with the greater increase in entropy? Explain your reasoning. (a) 2 H2O(,) : 2 H2(g) 1 O2(g) (b) C(s) 1 O2(g) : CO2(g) 10.43 Which reaction occurs with the greater increase in entropy? Explain your reasoning. (a) 2 NO(g) : N2(g) 1 O2(g) (b) Br2(g) 1 Cl2(g) : 2 BrCl(g) 10.44 Methanol is burned as fuel in some race cars. This makes it clear that the reaction is spontaneous once methanol is ignited. Yet the entropy change for the reaction 2 CH3OH(,) 1 3 O2(g) : 2 CO2(g) 1 4 H2O(,) is negative. Why doesn’t this violate the second law of thermodynamics? 10.45 Limestone is predominantly CaCO3, which can undergo the reaction CaCO3(s) : CaO(s) 1 CO2(g). We know from experience that this reaction is not spontaneous, yet DS for the reaction is positive. How can the second law of thermodynamics explain that this reaction is NOT spontaneous?


341

Describe what is happening in each beaker (a) on the molecular level and (b) in terms of the second law of thermodynamics.

The Third Law of Thermodynamics 10.46 Suppose that you find out that a system has an absolute entropy of zero. What else can you conclude about that system? 10.47

■ Use tabulated thermodynamic data to calculate the standard entropy change of each of the reactions listed below.

(a) Fe(s) 1 2 HCl(g) : FeCl2(s) 1 H2(g) (b) 3 NO2(g) 1 H2O(,) : 2 HNO3(,) 1 NO(g) (c) 2 K(s) 1 Cl2(g) : 2 KCl(s) (d) Cl2(g) 1 2 NO(g) : 2 NOCl(g) (e) SiCl4(g) : Si(s) 1 2 Cl2(g) 10.48 If you scan the values for S° in Appendix E, you will see that several aqueous ions have values that are less than zero. The third law of thermodynamics states that for a pure substance the entropy goes to zero only at 0 K. Use your understanding of the solvation of ions in water to explain how a negative value of S° can arise for aqueous species. 10.49 Calculate DS° for the dissolution of magnesium chloride: MgCl2(s) : Mg2+(aq) 1 2 Cl2(aq). Use your understanding of the solvation of ions at the molecular level to explain the sign of DS°. 10.50 Calculate the standard entropy change for the reaction CO 2(g) 1 2 H 2O(,) : CH 4(g) 1 2 O 2(g). What does the sign of DS° say about the spontaneity of this reaction? 10.51 Through photosynthesis, plants build molecules of sugar containing several carbon atoms from carbon dioxide. In the process, entropy is decreased. The reaction of CO2 with formic acid to form oxalic acid provides a simple example of a reaction in which the number of carbon atoms in a compound increases:

Gibbs Free Energy 10.56 Describe why it is easier to use DG to determine the spontaneity of a process rather than DSu. 10.57 Under what conditions does DG allow us to predict whether a process is spontaneous? 10.58 There is another free energy state function, the Helmholtz free energy (F), defined as F 5 E 2 TS. Comparing this to the definition of G, we see that internal energy has replaced enthalpy in the definition. Under what conditions would this free energy tell us whether or not a process is spontaneous? 10.59

Calculate DG° at 45°C for reactions for which

(a) DH° 5 293 kJ; DS° 5 2695 J/K (b) DH° 5 21137 kJ; DS° 5 0.496 kJ/K (c) DH° 5 286.6 kJ; DS° 5 2382 J/K 10.60

■ Discuss the effect of temperature change on the spontaneity of the following reactions at 1 atm.

(a) Al2O3(s) 12 Fe(s) : 2 Al(s) 1 Fe2O3(s) DH° 5 1851.5 kJ; DS° 5 138.5 J/K

CO2(aq) 1 HCOOH(aq) : H2C2O4(aq)

(b) N2H4(,) : N2(g) 1 2 H2(g) DH° 5 250.6 kJ; 1 (c) SO3(g) : SO2(g) 1 —— O2(g) 2 DH° 5 98.9 kJ;

(a) Calculate the standard entropy change for this reaction and discuss the sign of DS°. (b) How do plants carry out reactions that increase the number of carbon atoms in a sugar, given the changes in entropy for reactions like this?

■

10.61

■

DS° 5 0.3315 kJ/K

DS° 5 10.0939 kJ/K

The reaction CO2(g) 1 H2(g) : CO(g) 1 H2O(g)

10.52 Find websites describing two different attempts to reach the coldest temperature on record. What features do these experiments have in common?

is not spontaneous at room temperature but becomes spontaneous at a much higher temperature. What can you conclude from this about the signs of DH° and DS°, assuming that the enthalpy and entropy changes are not greatly affected by the temperature change? Explain your reasoning.

10.53 Look up the value of the standard entropy for the following molecules: CH4(g), C2H5OH(,), H2C2O4(s). Rank the compounds in order of increasing entropy and then explain why this ranking makes sense. 10.54 Look up the value of the standard entropy for the following molecules: SiO2(s), NH3(g), C2H6(g). Rank the compounds in order of increasing entropy and then explain why this ranking makes sense.

10.62 Why is the free energy change of a system equal to the maximum work rather than just the work?

10.55 A beaker of water at 40°C (on the left in the drawing) and a beaker of ice water at 0°C are placed side by side in an insulated container. After some time has passed, the temperature of the water in the beaker on the left is 30°C and the temperature of the ice water is still 0°C.

10.64

342

10.63 Distinguish between a reversible and an irreversible process.


For the reaction NO(g) 1 NO 2(g) : N 2O 3(g), use tabulated thermodynamic data to calculate DH° and DS°. Then use those values to answer the following questions.

■

(a) Is this reaction spontaneous at 25°C? Explain your answer.

■


(b) If the reaction is not spontaneous at 25°C, will it become spontaneous at higher temperatures or lower temperatures? (c) To show that your prediction is accurate, choose a temperature that corresponds to your prediction in part (b) and calculate DG. (Assume that both enthalpy and entropy are independent of temperature.)

(c) 2 K(s) 1 Cl2(g) : 2 KCl(s) (d) Cl2(g) 1 2 NO(g) : 2 NOCl(g) (e) SiCl4(g) : Si(s) 1 2 Cl2(g) 10.70

(a) Mg3N2(s) 1 6 H2O(,) : 2 NH3(g) 1 3 Mg(OH)2(s) (b) 4 CH3NH2(g) 1 9 O2(g) : 4 CO2(g) 1 10 H2O(,) 1 2 N2(g)

10.65 The combustion of acetylene is used in welder’s torches because it produces a very hot flame. 5 C2H2(g) 1 —— O2(g) : 2 CO2(g) 1 H2O(g) 2 DH° 5 21255.5 kJ (a) Use data in Appendix E to calculate DS° for this reaction. (b) Calculate DG° and show that the reaction is spontaneous at 25°C. (c) Is there any temperature range in which this reaction is not spontaneous? (d) Do you think you could use Equation 10.4 to calculate such a temperature range reliably? Explain your answer. 10.66 Natural gas (methane) is being used in experimental vehicles as a clean-burning fuel. (a) Write the equation for the combustion of CH4(g), assuming that all reactants and products are in the gas phase. (b) Use data from Appendix E to calculate DS° for this reaction. (c) Calculate DG° and show that the reaction is spontaneous at 25°C. 10.67 Silicon forms a series of compounds that is analogous to alkanes and has the general formula SinH2n12. The first of these compounds is silane, SiH4, which is used in the electronics industry to produce thin ultrapure silicon films. SiH4(g) is somewhat difficult to work with because it is pyrophoric at room temperature—meaning that it bursts into flame spontaneously. (a) Write an equation for the combustion of SiH4(g). (The reaction is analogous to hydrocarbon combustion, and SiO2 is a solid under standard conditions. Assume the water produced will be a gas.) (b) Use the data from Appendix E to calculate DS° for this reaction. (c) Calculate DG° and show that the reaction is spontaneous at 25°C. (d) Compare DG° for this reaction to the combustion of methane. (See the previous problem.) Are the reactions in these two exercises enthalpy or entropy driven? Explain. Free Energy and Chemical Reactions 10.68 Explain why DGf° of O2(g) is zero. 10.69

Using tabulated thermodynamic data, calculate DG° for these reactions.

■

(a) Fe(s) 1 2 HCl(g) : FeCl2(s) 1 H2(g) (b) 3 NO2(g) 1 H2O(,) : 2 HNO3(,) 1 NO(g)

■


Using tabulated thermodynamic data, calculate DG° for these reactions.

■

(c) Fe3O4(s) 1 4 CO(g) : 3 Fe(s) 1 4 CO2(g) (d) P4O6(s) 1 6 H2O(,) : 4 H3PO3(aq) 10.71 Calculate DG° for the dissolution of both sodium chloride and silver chloride using data from Appendix E. Explain how the values you obtain relate to the solubility rules for these substances. 10.72 Phosphorus exists in multiple solid phases, including two known as red phosphorus and white phosphorus. Consider the phase transition between white and red phosphorus: P4(s, white) : P4(s, red). Use data from Appendix E to determine which form of phosphorus is more stable at 25°C. Is this form of the solid more stable at all temperatures? If not, what temperatures are needed to make the other form more stable? 10.73

■ The normal melting point of benzene, C6H6, is 5.5°C. For the process of melting, what is the sign of each of the following? (a) DH°, (b) DS°, (c) DG° at 5.5°C, (d) DG° at 0.0°C, (e) DG° at 25.0°C

10.74 Calculate DG° for the complete combustion of one mole of the following fossil fuels: methane (CH 4), ethane (C2H6), propane (C3H8), and n-butane (C4H10). Identify any trends that are apparent from these calculations. 10.75

Estimate the temperature range over which each of the following reactions is spontaneous.

■

(a) 2 Al(s) 1 3 Cl2(g) : 2 AlCl3(s) (b) 2 NOCl(g) : 2 NO(g) 1 Cl2(g) (c) 4 NO(g) 1 6 H2O(g) : 4 NH3(g) 1 5 O2(g) (d) 2 PH3(g) : 3 H2(g) 1 2 P(g) 10.76

■ Recall that incomplete combustion of fossil fuels occurs when too little oxygen is present and results in the production of carbon monoxide rather than carbon dioxide. Water is the other product in each case.

(a) Write balanced chemical equations for the complete and incomplete combustion of propane. (b) Using these equations, predict which will have the larger change in entropy. (c) Use tabulated thermodynamic data to calculate DG° for each reaction. (d) Based on these results, predict the sign and value of DG° for the combustion of carbon monoxide to form carbon dioxide. INSIGHT INTO the Economics of Recycling 10.77 During polymerization, the system usually becomes more ordered as monomers link together. Could an endothermic polymerization reaction ever occur spontaneously? Explain.


343

10.78 If a particular polymerization reaction happened to be endothermic, how could it ever take place? Where would the increase in the entropy of the universe have to arise?

Substance

10.79 When polymers are recycled, the ends of the long-chain polymer molecules tend to break off, and this process eventually results in a degradation of physical properties, rendering the recycled polymer unusable. Explain why the breaking off of the ends of the polymer molecules is favorable from the standpoint of the entropy of the system.

10.85 Iodine is not very soluble in water, but it dissolves readily in a solution containing iodide ions by the following reaction: I2(aq) 1 I2(aq) : I32(aq) The following graph shows the results of a study of the temperature dependence of DG° for this reaction. (The solid line is a best fit to the actual data points.) Notice that the quantity on the y axis is DG°/T, not just DG°. Additional data relevant to this reaction are also given in the table below the graph.

0

40.83

Ga(,)

5.578

0.0888

59.25

Ga(g)

271.96

–0.050

S° ( J K21mol21)

0

233.76

169.03

186.2

274.8

(e) Suggest a reason why these two steps would need to be carried out separately.

DS° 5 ?

Ga(s)

197.674

(d) At what temperatures is reaction #1 spontaneous?

–0.052 ΔG°/T (kJ/K)

DGf° (kJ mol21)

126.8

(c) Calculate S° for O2(g).

10.83 Gallium metal has a melting point of 29.8°C. Use the information below to calculate the boiling point of gallium in °C. DHf° (kJ mol21)

2166.4

(b) Calculate DGf° for CO(g).

Calculate the entropy change, DS°, for the vaporization of ethanol, C2H5OH, at the boiling point of 78.3°C. The heat of vaporization of the alcohol is 39.3 kJ/mol.

Substance

2238.7

S° ( J K21mol21)

(a) Calculate DH° for reaction #1.

■

C2H5OH(,) : C2H5OH(g)

2110.5

CH3OH(,)

DGf° (kJ mol21)

NOTE: You should be able to work this problem without using any additional tabulated data.

Additional Problems

10.82

CO(g) CH4(g)

10.80 The recycling of polymers represents only one industrial process that allows creating order in one location by creating greater disorder at some other location, often at a power plant. List three other industrial processes that must create disorder in the surroundings to generate the desired material. 10.81 Diethyl ether is a liquid at normal temperature and pressure, and it boils at 35°C. Given that DH is 26.0 kJ/mol for the vaporization of diethyl ether, find its molar entropy change for vaporization.

DHf° (kJ mol21)

–0.054

–0.056


–0.058

For a metal, gallium has a very low melting point.

10.84

■ Methane can be produced from CO and H2. The process might be done in two steps, as shown below, with each step carried out in a separate reaction vessel within the production plant.

Reaction #1 Reaction #2

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CO(g) 1 2 H2(g) : CH3OH(,) DS° 5 2332 J/K 1 CH3OH(,) : CH4(g) 1 —— O2(g) 2 DS° 5 1162 J/K


–0.060

0.0030 0.0031 0.0032 0.0033 0.0034 0.0035 0.0036 0.0037 1/T (1/K)

DGf° (k J mol21)

Substance

I2(aq)

16.37

2

I (aq)

251.57

2

I3 (aq)

251.4

(a) Calculate DG° for this reaction at 298 K. (DO NOT read this value off the graph. Use the data given to calculate a more accurate value.) (b) Determine DH° for this reaction. Assume that DH° is independent of T. (HINT: You will need to use the graph provided to find DH°. It may help if you realize that the graph is a straight line and then try to write an equation for that line.)

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10.86

10.87

The enthalpy of vaporization for water is 40.65 kJ mol21. As a design engineer for a project in a desert climate, you are exploring the option of using evaporative cooling. (a) If the air has an average volumetric heat capacity of 0.00130 J cm 23 K 21, what is the minimum mass of water that would need to evaporate in order to cool a 5 m 3 5 m room with a 3 m ceiling by 5°F using this method? (b) Is this a spontaneous or nonspontaneous process? ■

Determine whether each of the following statements is true or false. If false, modify to make the statement true.

■

(a) An exothermic reaction is spontaneous. (b) When DG° is positive, the reaction cannot occur under any conditions. (c) DS° is positive for a reaction in which there is an increase in the number of moles. (d) If DH° and DS° are both negative, DG° will be negative. 10.88

■ Nickel metal reacts with carbon monoxide to form tetracarbonyl nickel, Ni(CO)4.

Ni(s) 1 4 CO(g) : Ni(CO)4(g) This reaction is exploited in the Mond process in order to separate pure nickel from other metals. The reaction above separates nickel from impurities by dissolving it into the gas phase. Conditions are then changed so that the reaction runs in the opposite direction to recover the purified metal. (a) Predict the signs of DH° and DS° for the reaction as written above. (Note that bonds are formed but none are broken.) (b) Use tabulated thermodynamic data to calculate DH°, DS°, and DG° for the reaction. (c) Find the range of temperatures at which this reaction is spontaneous in the forward direction. 10.89

Polyethylene has a heat capacity of 2.3027 J g21 °C21. You need to decide if 1.0 ounce of polyethylene can be used to package a material that will be releasing heat when in use. Consumer safety specifications indicate that the maximum allowable temperature for the polyethylene is 45°C; it can be assumed that the plastic is initially at room temperature. (a) What temperature will the polyethylene reach if the product generates 1500 J of heat and all of this energy is absorbed by the plastic package? (b) Is this a realistic estimate of the temperature that the polyethylene packaging would reach? Explain your answer. (c) What is the enthalpy change of the polyethylene? (d) Estimate the entropy change of the polyethylene. (You will need to assume that the temperature of the plastic is constant.)

■

10.90 A key component in many chemical engineering designs is the separation of mixtures of chemicals. (a) What happens to the entropy of the system when a chemical mixture is separated? (b) Are designs for chemical separation more likely to rely on spontaneous or nonspontaneous processes?

■


10.91

The reaction shown below is involved in the refining of iron. (The table that follows provides all of the thermodynamic data you should need for this problem.) ■

2 Fe2O3(s) 1 3 C(s, graphite) : 4 Fe(s) 1 3 CO2(g) Compound

Fe2O3(s)

DHf° S° DGf° (kJ mol21) (J mol21 K21) (kJ mol21)

2824.2

C(s, graphite)

0

Fe(s)

0

CO2(g)

?

2742.2

5.740

0

27.3

2393.5

213.6

0 2394.4

(a) Find DH° for the reaction. (b) DS° for the reaction above is 557.98 J/K. Find S° for Fe2O3(s). (c) Calculate DG° for the reaction at the standard temperature of 298 K. (There are two ways that you could do this.) (d) At what temperatures would this reaction be spontaneous?

FOCUS ON PROBLEM SOLVING EXERCISES 10.92 Suppose that you need to know the heat of formation of 2-pentene, C5H10, but the tables you have do not provide the value. You have a sample of the chemical. What could you do to determine the heat of formation? 10.93 Suppose that you need to know the heat of formation of cyclohexane, C6H12, but the tables you have do not provide the value. You have a sample of the chemical. What could you do to determine the heat of formation? 10.94 You have a table of thermodynamic variables that includes heats of formation and standard entropies but not free energies of formation. How could you use the information you have to estimate the free energy of formation of a substance that is listed in your table? 10.95 You have a table of thermodynamic variables that includes standard entropies and free energies of formation but not heats of formation. How could you use the information you have to estimate the heat of formation of a substance that is listed in your table? 10.96 From a series of hydrocarbons containing only single carbon2carbon bonds, how could you identify a trend for the heat of formation for successive carbon–carbon bond formation? 10.97 Thermodynamics provides a way to interpret everyday occurrences. If you live in northern climates, one common experience is that during early winter, snow falls but then melts when it hits the ground. Both the formation and the melting happen spontaneously. How can thermodynamics explain both of these seemingly opposed events? 10.98 Suppose that you are designing a chemical reactor and you need to know the heat of vaporization for a solvent. You don’t have any thermodynamic tables handy, but you know the boiling point of the liquid and your CAD program can calculate third-law entropies. How would you determine the heat of vaporization?


345

Cumulative Problems 10.99 Fluorine reacts with liquid water to form gaseous hydrogen fluoride and oxygen. (a) Write a balanced chemical equation for this reaction. (b) Use tabulated data to determine the free energy change for the reaction and comment on its spontaneity. (c) Use tabulated data to calculate the enthalpy change of the reaction. (d) Determine how much heat flows and in what direction when 34.5 g of fluorine gas is bubbled through excess water. 10.100 Ammonia can react with oxygen gas to form nitrogen dioxide and water. (a) Write a balanced chemical equation for this reaction. (b) Use tabulated data to determine the free energy change for the reaction and comment on its spontaneity. (c) Use tabulated data to calculate the enthalpy change of the reaction. (d) Determine how much heat flows and in what direction when 11.4 g of ammonia gas is burned in excess oxygen. 10.101 Consider the following thermodynamic data for oxides of manganese.

346

DHf° (kJ mol21)

DGf° (kJ mol21)

S° ( J K21mol21)

MnO

2385.2

2362.9

59.71

MnO2

2520.0

2465.2

53.05

Mn2O3

2959.0

2881.2

Substance

Mn3O4

21388

21283

110.5 155.6

(a) What is the correct chemical nomenclature for each of the first three oxides? (b) Write and balance chemical equations for the conversion of each of these oxides into Mn3O4. (c) Based on the free energy changes of these reactions, which oxide is the most stable at room temperature? 10.102 (a) When a chemical bond forms, what happens to the entropy of the system? (b) Thermodynamically, what allows for any bond formation to occur? (c) What do your answers to parts (a) and (b) suggest must be true about the formation of chemical bonds for the octet rule to hold?


■


Chemical Kinetics

11 OUTLINE 11.1 INSIGHT INTO Ozone Depletion 11.2 Rates of Chemical Reactions 11.3 Rate Laws and the Concentration Dependence of Rates 11.4 Integrated Rate Laws 11.5 Temperature and Kinetics 11.6 Reaction Mechanisms 11.7 Catalysis 11.8 INSIGHT INTO Tropospheric Ozone

Chlorine in the stratosphere has been implicated as a major cause of ozone depletion. Elevated chlorine concentration in the colored regions in the first column correlates strongly with ozone loss, as represented by the blue regions in the center column. European Space Agency

W

hen we considered thermodynamics, we concentrated on what should eventually happen to a chemical system. But we have not yet given any thought to how long it might take to reach this eventual outcome or exactly how a collection of reactants might be transformed into products. Fundamentally, chemistry is about change. So if we want to understand chemistry, we will need to examine both the process of change and the speed with which it occurs. On the Fourth of July, it is exhilarating to watch the explosive oxidation reactions of fireworks, but a holiday event to watch iron rust would probably not draw a crowd! What factors determine whether a reaction proceeds with explosive speed or at an imperceptible rate? Could we learn to manipulate these factors to force a reaction to occur at the speed we choose? The area of chemistry dealing with these questions of “How?” or “How fast?” is called chemical kinetics, and we begin with a glimpse into stratospheric ozone and its depletion.

Chapter Objectives After mastering this chapter, you should be able to ❚ explain the role of chemical kinetics in the formation and destruction of ozone in the atmosphere. ❚ define the rate of a chemical reaction and express the rate in terms of the concentrations of individual reactants or products. ❚ use the method of initial rates to determine rate laws from experimental data. ❚ use graphical methods to determine rate laws from experimental data.

Online homework for this chapter may be assigned in OWL. 347

❚ explain the difference between elementary reactions and multistep reactions. ❚ find the rate law predicted for a particular reaction mechanism. ❚ use a molecular perspective to explain the significance of the terms in the Arrhenius equation. ❚ calculate the activation energy for a reaction from experimental data. ❚ explain the role of a catalyst in the design of practical chemical reactions.

INSIGHT INTO

11.1 300

Altitude (km)

200

Thermosphere

100 80

Mesosphere

60 40 20

Stratosphere Troposphere Earth

Figure 11.1 ❚ Several layers of the atmosphere are identified in the figure. The ozone layer is in the stratosphere at an altitude of about 30 km.

Sydney Chapman was a mathematician. He devised the idea of the Chapman cycle while working on a mathematical theory of gases.

Other measurements make clear that the depletion of ozone over polar areas of the planet is a seasonal process.

348

Ozone Depletion

Walk outside right after a summer thunderstorm and you will probably notice a pleasantly fresh smell. The source of that enjoyable scent is ozone (O3), the less common allotrope of oxygen. Though its pungent odor can be pleasing in very low concentration, it becomes unpleasant and even toxic in concentrations greater than about one part per million, causing headache and difficult breathing. In a thunderstorm, energy from lightning helps drive reactions that produce ozone from ordinary oxygen. But ozone can also be formed through reactions of various compounds in the exhaust gases from automobiles and industrial processes. Weather forecasts for many urban areas now include “ozone alerts” on days when the ozone level in the air is expected to rise above about 0.1 ppm. Ozone is a major air pollutant in the troposphere (the atmosphere at ground level—see Figure 11.1), and yet a major topic in the news is concern about the depletion of the stratospheric ozone layer. Why is ozone a problem in the troposphere but a necessity in the stratosphere? What is the chemistry behind these issues? The smell of ozone after a thunderstorm disappears quickly, partly because it diffuses through the air (and because your olfactory nerves dull) but also because ozone is unstable and decomposes to form ordinary oxygen. The decomposition reaction has a deceptively simple overall equation: 2 O3 : 3 O2. That ozone is reactive and cannot exist for long at the earth’s surface suggests two important facts. First, O2 is the more stable of the two allotropes of oxygen. So thermodynamics must favor the decomposition of ozone. Second, for the ozone layer to exist, the upper atmosphere must have some conditions that allow the formation of the less stable allotrope, O3, in significant concentration. The accumulation of ozone in the stratosphere is governed by the way these conditions influence the rates of chemical reactions and therefore provides our initial insight into the role of kinetics in chemistry. In 1903, the British scientist, Sydney Chapman, first explained the chemistry of the formation of ozone in the upper atmosphere when he proposed what is now known as the Chapman cycle. Although Chapman’s original proposal was quite speculative, modern measurements of a wide array of quantities support his hypothesis. The Chapman cycle begins and ends with diatomic oxygen, the more stable allotrope, as shown in Figure 11.2. The first step is the photochemical dissociation of O2 to form oxygen atoms, which may then react with O2 molecules to form ozone, O3. Mechanisms also exist for ozone decomposition, so that, even in the stratosphere, ozone is not stable. So why does the ozone layer exist? To answer this, we need to consider the rates at which ozone is produced and consumed. If ozone is produced faster than it is consumed, its concentration will increase. But if ozone is destroyed faster than it is formed, the ozone level will decrease. The existence of the ozone layer, then, depends on this balance between the rates at which ozone is produced and destroyed. To understand ozone depletion, we will need to understand reaction rates. The ozone layer would probably not make headlines were it not at risk. During the past few decades, scientists have been measuring a decrease in ozone concentration that returns during the early spring months, particularly over Antarctica and (less dramatically) over North America. Evidence of this effect is presented graphically in Figure 11.3. The decrease in stratospheric ozone concentration from that of 30 years ago is what is commonly referred to as the ozone hole.

Chapter 11 Chemical Kinetics

Figure 11.2 ❚ The Chapman

1 The Chapman cycle begins with a UV photon dissociating a diatomic oxygen molecule.

UV

cycle for the formation and destruction of ozone in the stratosphere.

2 An oxygen atom that results from this dissociation may collide with another diatomic oxygen to form ozone.

We can draw two equivalent resonance structures for ozone. O ⫹ O3 4 Another fate of ozone is that it may collide with an oxygen atom and reform the original starting material — diatomic oxygen. There are, however, relatively few oxygen atoms around, so such collisions are less common than the other reactions shown.

O

O

O

O

O

O

UV 3 One fate of the ozone formed is that it may absorb a different UV photon. Doing so results in the dissociation of the ozone. This reaction is the reason ozone shields us from excess UV radiation.

Figure 11.3 ❚ The data in this graph clearly show a decrease in the overall levels of atmospheric ozone from the late seventies through the early nineties. The increase in ozone levels in recent years is due mainly to the impact of the Montreal Protocol, which curtailed the use of ozone damaging substances. Source: WMO [World Meteorological Organization] Scientific Assessment of Ozone Depletion: 2006, Global Ozone Research and Monitoring Project, Report No. 50, Geneva, 2007.)

Global Total Ozone Change Changes from 1964 –1980 average

2

Ozone Change (%)

0

–2

–4

Average

–6

Range of observations

–8 1965

1975

1985 Year

1995

2005

11.1 Ozone Depletion

349

How does the ozone hole arise? We must consider the role of species other than the allotropes of oxygen in the Chapman cycle. There is evidence that chlorine and bromine in the stratosphere lead to a decrease in the amount of ozone present. How does this occur? What factors influence whether or not the decrease occurs? And what might be done to combat the loss of ozone? These questions all point to aspects of chemical kinetics that we will address in this chapter. To begin this study, we focus first on the concepts of rate of reaction and the ways in which reaction rate can be measured.

11.2

Rates of Chemical Reactions

We suggested that the rates of different chemical reactions govern ozone levels in the atmosphere, and this seems sensible. If ozone—or anything else—is produced faster than it is consumed, it will accumulate. To discuss these ideas quantitatively, we must first raise two fundamental issues in chemical kinetics. How do we define the rate of the reaction? And how can we measure it?

Concept of Rate and Rates of Reaction Aside from our interest in the rate of destruction of ozone, we have some general experience with rates. If we travel 55 miles from home and it takes us an hour to go that distance, we say that our average speed is 55 miles per hour. distance traveled Average speed 5 — time elapsed In this ratio, there is an indication of progress (distance traveled) related to the time required to make that progress. Because the distance measured was from zero miles (at home) to 55 miles from home, the change in distance from home can be symbolized by Dd, where D is the symbol for change. The time changed from zero time (at home) to one hour later at our destination. So the time elapsed can be represented by Dt. Dd Average speed 5 — Dt How do we translate this ratio into one that is meaningful for chemical reactions? We need a ratio of the progress of the reaction to the time it takes to make that progress. In a chemical reaction, the identities of the chemicals change, rather than the distance traveled. So, to define the chemical rate, we must measure the change in chemical content that occurs during the reaction. Usually we do this by measuring concentrations. Thus the reaction rate is the ratio of the change in concentration to the elapsed time. change in concentration Rate 5 —— elapsed time The timescales of interest in chemical kinetics can vary over a very wide range, from picoseconds to years.

(11.1)

Notice the units implied by this ratio. Concentration generally is measured in molarity or mol L21 and time in seconds, so rate units are usually mol L21 s21. We can express this ratio in a mathematically compact form by using square brackets, [ ], to designate a concentration or molarity. D[substance] Rate 5 — Dt Using our example of ozone decomposition, now, we can begin to write an expression for the rate of reaction D[O3] Rate 5 — Dt

350


Stoichiometry and Rate The overall reaction for the destruction of ozone was mentioned in the opening “Insight” section. 2 O3 : 3 O2 As we saw above, the rate of this reaction can be given by the rate at which ozone decomposes. D[O3] Rate 5 — Dt But suppose that we had monitored the increase in O2 concentration over time, instead of the decrease in ozone concentration. The change in concentration we’d measure would be different, yet it must also be determined by the rate of the same reaction. As the reaction proceeds, oxygen concentration increases more quickly than ozone concentration decreases because three oxygen molecules form for every two ozone molecules destroyed. Which rate should we measure, the rate of O2 production or the rate of disappearance of O3? The answer is either one, but to do so consistently, we must take into account the stoichiometry of the reaction. If we are measuring the increase in the concentration of the product, the rate of a reaction is naturally a positive number. When we observe a decrease in the concentration of a reactant, we must include a minus sign in the rate statement to obtain a positive value for the rate. Thus rate might be defined as either D[product] Rate 5 — Dt or D[reactant] Rate 5 2 — Dt But this still does not account for the fact that the number of O2 molecules produced in a given time is greater than the number of O3 molecules consumed. To account for this, we also include the stoichiometric coefficient in the denominator of the rate expression: D[product] Rate 5 — nprodDt

(11.2.a)

D[reactant] Rate 5 2 — nreactDt

(11.2.b)

or

Here nprod is the stoichiometric coefficient of the product being measured, and nreact is the coefficient of the reactant. Accounting for stoichiometry in this way ensures that the same reaction rate is obtained using either the rate of disappearance of reactants or the rate of appearance of products.


11.1

The conversion of ozone to oxygen, 2 O3 : 3 O2, was studied in an experiment, and the rate of ozone consumption was measured as 2.5 3 1025 mol L21 s21. What was the rate of O2 production in this experiment?

Strategy Use the stoichiometry of the reaction to relate the respective rate expressions. Because the rate of the reaction is the same whichever quantity is measured, we can set the expressions for the rate of oxygen appearing and ozone disappearing equal to each other and solve for the unknown quantity.

11.2 Rates of Chemical Reactions

351

Solution Begin by writing the expressions for the rate of reaction in terms of both O2 and O3 concentration. D[O3] D[O2] Rate 5 2 — 5 — 2Dt 3Dt Remember that the change in concentration of the reactant is a negative quantity and D[O3] substitute 22.5 3 1025 mol L21 s21 for —. Dt 25 D[O2] 2.5 3 10 mol L21 s21 —— 5 — 2 3Dt Solve for the rate of O2 formation. Δ[O2] Δt

3 (2.5 3 1025 mol L21 s21) 2

— 5 ——

5 3.8 3 1025 mol L21 s21

Discussion For some students, the most troubling part of this mathematical description of rate is the minus sign when reactant concentrations are used. By recalling our standard definition for D, however, we can see the need for this sign more explicitly. In this case, we are interested in D[O3], which is [O3]later 2 [O3]initial. The initial concentration of a reactant is larger than its concentration at some later time, so D[O3] is negative. We need the minus sign in the equation so that the rate is not a negative number.

Check Your Understanding Like ozone, dinitrogen pentoxide decomposes

to form O2 (along with nitrogen dioxide) by the reaction 2 N2O5 : 4 NO2 1 O2. If the rate of disappearance of N2O5 is 4.0 3 1026 mol L21 s21, what is the rate of appearance of each product?

Average Rate and Instantaneous Rate

Figure 11.4 ❚ When a candle burns in a closed container, the flame will diminish and eventually go out. As the amount of oxygen present decreases, the rate of combustion will also decrease. Eventually, the rate of combustion is no longer sufficient to sustain the flame even though there is still some oxygen present in the vessel.

352


Photos: Thomas A. Holme

Let’s think about the simple experiment like the one shown in Figure 11.4. Take a glass and place a candle inside. Light the candle, let it burn for a few moments, and cover the glass with a saucer. Observe the flame. What happens? The brightly burning flame gradually decreases in size until it goes out because the concentration of one reactant, oxygen, has decreased. During this experiment, the rate of the combustion reaction decreases until the reaction stops. What is the rate of the reaction? Is it the rate at which oxygen is consumed in the brightly burning flame or in the nearly extinguished one? This simple experiment shows that the rate we observe depends on when or how we make our observations.

Figure 11.5 ❚ The distinction between the instantaneous rate of reaction and the average rate measured over some period of time is illustrated here. The slope of the green line to the left gives the instantaneous rate at time t1. The slope of the blue line to the right gives the average rate for the period from time t1 to time t2.

The instantaneous rate is the slope of a tangent line at a single point on the curve.

Concentration

The avergage rate is the slope of a line between two points.

t1

Time

t2

Figure 11.5 shows a graph of the concentration of ozone as a function of time in a laboratory experiment. Looking closely at this figure, we can see the slight difference between the average reaction rate and the instantaneous reaction rate. The difference between these two methods of defining rate lies in the amount of time needed to make the observation. For the average rate, two concentrations are measured at times separated by a finite difference, and the slope of the line between them gives the rate. The instantaneous rate refers to the rate at a single moment, and it is given by the slope of a line tangent to the curve defined by the change in concentration versus time. As Figure 11.5 shows, these two slopes can differ. (Returning to the speed analogy, your instantaneous speed at any moment during a long drive might be very different from your average speed during the same trip. You may accelerate to pass another car or come to a stop in traffic.) In most cases, we prefer to work with instantaneous rates in kinetics. One commonly measured rate is the initial rate: the instantaneous rate of the reaction just as it begins. In Figure 11.5, the tangent line and the concentration curve coincide for at least a brief period at the beginning of the reaction, so the initial instantaneous rate is often relatively easy to measure. To determine rates of reaction in complicated systems such as the stratosphere, we typically try to measure the rates of the same reactions independently in the laboratory. Thus to consider the reactions of the Chapman cycle, a chemist would have to devise an experiment that measures the rate of reaction of some very reactive species, such as oxygen atoms. The reactive species is first created (maybe by using a pulse of laser light to break the bond in O2), and then its rapid consumption is monitored. This can be a difficult measurement because the reaction may proceed extremely quickly. Reactive species such as oxygen atoms react with most chemicals in the reaction container, so great care must be taken to isolate the reacting species to ensure that only the reaction of interest can occur. Once concentrations of reacting species are measured, an equation can be formulated to describe how the rate depends on these concentrations.

11.3

If you are familiar with calculus, you’ll realize that the instantaneous rate is the derivative of concentration with respect to time.

Rate Laws and the Concentration Dependence of Rates

We saw in Figure 11.4 that as O2 is consumed, the rate of combustion decreases. The rate of a chemical reaction depends on a number of factors. One of these factors, the concentration of the reacting species, can help us understand why there is an ozone

11.3


353

layer in the upper atmosphere, despite the fact that O2 is the more stable of the two oxygen allotropes. To approach this issue for the stratosphere, let’s first investigate how reaction rates measured in a laboratory depend on concentration.

The Rate Law Observations of many chemical reactions show that the dependence of reaction rate on concentration often follows relatively simple mathematical relationships. This behavior can be summarized in a mathematical equation known as the rate law. There are two useful forms of the rate law, and we begin with the differential rate law (the name and the relationship are derived from calculus). For a reaction between substances X and Y, the rate of reaction can usually be described by an equation of the form Rate 5 k[X]m[Y]n

(11.3)

where k is a constant called the rate constant, [X] and [Y] represent the reactant concentrations, and m and n are typically either integers or half integers. The actual values of the exponents m and n must be measured experimentally. (In some cases, these values can be related to the stoichiometric coefficients of the reaction, but not always!) The experimentally determined values of the exponents, in this case m and n, are referred to as the order of the reaction. For example, if m 5 1, the reaction is first order with respect to reactant X; if m 5 2, the reaction is second order with respect to X. It is unusual for any exponent in the rate law to have a value greater than two. When a rate law depends on more than one reactant concentration, we can distinguish between the overall order of the reaction and the order with respect to individual reactants. Thus a rate law such as Rate 5 k[A][B] is second order overall but first order with respect to reactant A and first order with respect to reactant B. (Each exponent is understood to be one.)

E XAM PL E PRO BLEM

11.2

In the following rate laws, determine the orders with respect to each substance and the overall order of the reaction. (a) Rate 5 k[A]2[B], (b) Rate 5 k[A][B]1/2

Strategy The order with respect to each individual species is the exponent, and the overall order is the sum of the individual orders.

Solution (a) The order with respect to A is 2 and with respect to B is 1. The overall order is 3. (b) The order with respect to A is 1 and with respect to B is 1/2. The overall order is 3/2.

Discussion It is important to remember that the order of the reaction is determined experimentally. Although this problem focuses on the mathematical detail of the definition of reaction order, we must keep in mind that reaction order and stoichiometry are not directly related.

Check Your Understanding The following reactions were studied in the laboratory, and their respective rate laws were determined. Find the reaction orders with respect to each reactant and the overall order in each of the following: (a) H2(g) 1 Br2(g) : 2 HBr; rate 5 k[H2][Br2] (b) 2 N2O5 : 4 NO2 1 O2; rate 5 k[N2O5]

354


The rate constant gives us valuable information about the kinetics of a chemical reaction. Perhaps most importantly, the magnitude of the rate constant tells us whether or not a reaction proceeds quickly. If the rate constant is small, the reaction is likely to proceed slowly. By contrast, a large rate constant indicates a rapid reaction. Rate constants for chemical reactions range over many orders of magnitude. We should also point out that the value of the rate constant depends on the temperature of the reaction. (We will consider temperature dependence further in Section 11.5, but it may take some time to get used to the idea of a “constant” that is a function of a variable.) The temperature dependence of the rate of a reaction is described in terms of the rate constant, as we’ll see in Section 11.5. The units of a rate constant depend on the overall order of the reaction. We have already established that rate has the units of mol L21 s21 when concentrations are given in mol L21. The units of the rate constant must be chosen so that units balance in the rate law. You should be able to verify that the units for the rate constant of a first-order reaction must be s21, whereas for second-order reactions, the rate constant has the units L mol21 s21. Not all rate laws are as simple as those mentioned thus far, all of which have included only concentrations of reactants. One reaction for the destruction of stratospheric O3, for example, has a concentration dependence on the product, O2. Increased O2 concentration, however, slows the reaction rather than increasing it. This observation requires a negative exponent for oxygen in the rate law. Thus the rate law for 2 O3 : 3 O2 is given by [O3]2 Rate 5 k — [O2] or Rate 5 k[O3]2[O2]21 It is also not uncommon to find orders of 1/2 or 21/2 in rate laws for some reactions.

Determination of the Rate Law As noted earlier, a rate law must be determined experimentally. There are two common ways to do this. One is to use a series of graphs to compare data to various possible rate laws, and that method is emphasized in Section 11.4. First we will look at the other approach, which is to measure the initial rate of the reaction while adjusting the concentrations of the various reactants. To see the logic needed to determine a rate law, let’s think about a simple reaction in which some substance we’ll label as A is the only reactant. The rate for such a reaction should be given by a rate constant times the concentration of A raised to some power. Rate 5 k[A]n Keep in mind that the order of reaction (n) is usually an integer and is rarely greater than two. That gives us three likely orders: 0, 1, and 2. In each of these cases, if the concentration of A is doubled, the rate will change in a simple and predictable way. 1. If n 5 0, doubling the concentration of A does not change the rate at all because any quantity raised to the 0 power is 1. 2. If n 5 1, doubling the concentration of A doubles the rate. 3. If n 5 2, doubling the concentration of A increases the rate by a factor of 4 (because 22 5 4). Other orders of reaction are also possible and can give more complicated relationships. But these three cases cover the great majority of the reactions we will encounter. Let’s determine the rate law for such a reaction involving only one reactant.

11.3


355

E XAM PL E PRO BLEM

11.3

Earlier in the chapter we mentioned the decomposition of N2O5: 2 N2O5(g) : 4 NO2(g) 1 O2(g) Consider the following data for the kinetics of this reaction.

Experiment

Initial [N2O5] (mol L21)

Initial Rate of Reaction (mol L21 s21)

1

3.0 3 1023

9.0 3 1027

2

9.0 3 1023

2.7 3 1026

Determine the rate law and rate constant for this reaction at the temperature of these experiments.

Strategy To establish the rate law, determine the order with respect to the reactant by noticing how the rate changes with changing concentration. Write the rate law and calculate k by substituting the concentration and rate from one of the experiments in the rate law.

Solution Looking at the data, we see that the initial concentration triples from the first to the second experiment. 3(3.0 3 1023) 5 9.0 3 1023 mol L21 And, the rate also triples: 3(9.0 3 1027) 5 2.7 3 1026 mol L21 s21 Because tripling the concentration triples the rate, the reaction must be first order with respect to N2O5. Rate 5 k[N2O5] Now use the values from either experiment to determine k. Using Experiment #1 from the table, 9.0 3 1027 mol L21 s21 5 k(3.0 3 1023 mol L21) Therefore, k 5 3.0 3 1024 s21.

Analyze Your Answer This is a problem where we have little initial intuition, but we can still take a quick look at the numbers and see if they are reasonable by focusing on the orders of magnitude. Looking at only the powers of 10, the rates measured and the concentrations used in the experiments have differences of either 103 or 104. Getting a value for k in this range, therefore, makes sense.

Discussion The absolute size of the initial concentration does not matter when we carry out this type of experiment. Neither does the choice to double or triple that initial concentration. We could equally well have chosen to cut the initial concentration in half and then measured the new rate. The logic used to determine the rate law remains the same. If we were doing the experiment, we would probably choose to change the concentration by a factor of 2 or 3 rather than by some arbitrary amount, but even that is not really necessary. It just makes it easier to recognize the expected effects for each possible order of reaction.

Check Your Understanding Would the result change if we used Experiment 2 to determine the rate constant?

356


Now let’s confront the situation in which two reactants are used, and we need to determine the order with respect to each of them. In this case, our rate law will have three unknowns: the rate constant and the two reaction orders. To determine values for these three parameters, we must carry out at least three experiments. (This follows from the idea that at least three equations are needed to solve for three unknowns.) When working with a system that depends on several variables, it is always a good idea to try to separate the influence of one variable from the others. In this case, we can do that by holding one reactant concentration constant while changing the other to determine its effect on the rate.


11. 4

The study of the kinetics of real systems can be complicated. For example, there are several ways by which O3 can be converted to O2. One such reaction is NO2 1 O3 : NO3 1 O2 Three experiments were run, and the following data obtained.

Experiment

Initial [NO2] (mol L21)

Initial [O3] (mol L21)

Initial Rate of Reaction (mol L21 s21)

1

2.3 3 1025

3.0 3 1025

1.0 3 1025

2

4.6 3 1025

3.0 3 1025

2.1 3 1025

3

4.6 3 1025

6.0 3 1025

4.2 3 1025

Determine the rate law and rate constant for this reaction.

Strategy To determine the order with respect to each of the reactants, look for pairs of experiments that differ only in the concentration of one reactant and compare their rates. Any difference in rate must be due to the effect of the reactant whose concentration changed. Write the rate law and use it to determine k.

Solution Order with respect to NO2 From Experiments 1 and 2, we see that, if we double [NO2] while holding [O3] constant, we double the rate (within experimental error). Thus the order in NO2 is 1. Order with respect to O3 From Experiments 2 and 3, we see that if we double [O3] while holding [NO2] constant, we double the rate. Thus the order in O3 must also be 1. The rate law is

Because rates are determined experimentally, the change in rate may not be an exact multiple of an integer.

Rate 5 k[NO2][O3] Evaluate k, using data from any of the three experiments. Here we will use Experiment 2. Rate 5 k[NO2][O3] 25

2.1 3 10

21 21

mol L

s

5 k(4.6 3 1025 mol L21)(3.0 3 1025 mol L21)

k 5 1.5 3 104 L mol21 s21

Analyze Your Answer Once again, we can take a quick look to see if the numbers

are reasonable by focusing on the powers of 10. We have two factors of 1025 (the concentrations) on the right-hand side of the equation and one factor of 1025 (the rate) on the left. Getting a value for k that roughly offsets one of the concentrations makes sense.

11.3


357

Discussion The key to solving this type of problem lies in recognizing which concentrations change and which stay the same. Sometimes when rates are reported using scientific notation, it is easy to miss a change in the exponent. These types of organizational details are the only added complication of this type of problem relative to Example Problem 11.3.

Check Your Understanding Why are the units for k in this example different from the units in Example Problem 11.3?

11.4 The differential and integrated forms of the rate law are both equally valid. The integrated rate law is useful if we want to consider concentrations as a function of time.

Integrated Rate Laws

The rate law determined by the method described in the previous section allows us to predict the rate of a reaction for a given set of concentrations. But because the concentrations of the reacting substances will change with time, the rate law does not let us easily predict the concentrations or rate at some later time. To do that, we need to formulate an equation that tells us explicitly how the concentrations will change as a function of time. This new equation, called the integrated rate law, can be derived from the rate law itself. (See MathConnections box below.) The reasoning behind the integrated rate law is this: If we know all of the concentrations in a reacting system at any given time, and we know how the reaction rate changes with time, then we should be able to predict what the concentrations will be at some later time. This is similar in concept to a physics problem where you might calculate the position of a projectile from its initial position and velocity. The form of the integrated rate law depends on the order of reaction. If we know the integrated rate laws for a few common reaction orders, we can use them as models for comparison with data. As we’ll see below, this provides a useful alternative way of determining the rate law for a reaction.

MathConnections You may have already learned about integrals in calculus. If so, you could probably guess from the name that the integrated rate law is found by integrating the rate law. Let’s take a look at how that’s done. (If you have not learned about integrals yet, don’t worry. You can work with the integrated rate law without knowing how to derive it.) The simplest case to start with is a zero-order reaction with a single reactant we’ll again call A. For such a case, the rate law is just Rate 5 k[A]0 5 k Now remember that the rate is the change in concentration over the change in time. We’ve been writing this in terms of macroscopic changes, using the ‘D’ notation. D[A] Rate 5 2 — 5 k Dt But if we simply let the changes become infinitesimal, we can convert from deltas to differentials. d[A] 2— 5 k dt We can rearrange this to put one of the differential terms on each side. d[A] 5 2kdt Now we can take the integral of both sides to eliminate the differentials. We will start at the beginning of the reaction, when t 5 0 and [A] 5 [A]0, and

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integrate from there to some later time when the corresponding values are just t and [A]t. [A]t

t

0

0

3[A] d[A] 5 2k3 dt Both of these integrals are simple to evaluate, and we get the following result: [A]t 2 [A]0 5 2kt This is the integrated rate law for the zero-order reaction. If you are familiar with integration, you should be able to follow the same procedure to derive the integrated rate laws for first-order and second-order kinetics.

Zero-Order Integrated Rate Law The simplest rate law is for a zero-order reaction.

Zero-order rate laws are relatively rare.

Rate 5 k[A]0 5 k In such instances, the rate of the reaction will not change as the reactants are consumed. The integrated rate law for this type of kinetics is given by [A]t 5 [A]0 2 kt

(11.4a)

Rearrange this equation slightly and compare it to the equation for a straight line with the form y 5 mx 1 b. [A]t 5 2kt 1 [A]0

(11.4b)

y 5 mx 1 b Because the forms of these expressions match, we can see that if we plot [A] (on the y axis) as a function of t (on the x axis), we will get a straight line. We can also see that the slope of that line (m) must be equal to 2k and the y intercept (b) must be equal to [A]0, the initial concentration of reactant A. Equation 11.4 provides us a model of the behavior expected for a system obeying a zero-order rate law. To test this model, we simply need to compare it with data for a particular reaction. So we could measure the concentration of reactant A as a function of time and then plot [A] versus t. If the plot is linear, we could conclude that we were studying a zero-order reaction. The catalytic destruction of N2O in the presence of gold is an example of this type of kinetics. A graphical analysis of the reaction is shown in Figure 11.6.

Catalysis and other catalytic reactions are discussed in Section 11.7.

Figure 11.6 ❚ The decomposition

[N2O] For a zero-order reaction, a plot of reactant concentration vs. time is linear. The slope of the line is –k.

of N2O over gold is a zero-order process, for which the reactant concentration decreases linearly as a function of time.

Time


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First-Order Integrated Rate Law For a first-order reaction with a single reactant, A, the integrated rate law is given by the equations [A]t ln — 52kt [A]0

(11.5)

[A]t 5 [A]0e2kt

(11.6)

or [A]0 represents the initial concentration of the reactant at time t 5 0. With these equations, if we know the rate constant and the initial concentration, we can predict the concentration at any subsequent time, as shown in Example Problem 11.5.

E XAM PL E PRO BLEM

11.5

The photodissociation of ozone by ultraviolet light in the upper atmosphere is a firstorder reaction with a rate constant of 1.0 3 1025 s21 at 10 km above the planet’s surface. O3 1 hn : O 1 O2 Consider a laboratory experiment in which a vessel of ozone is exposed to UV radiation at an intensity chosen to mimic the conditions at that altitude. If the initial O3 concentration is 5.0 mM, what will the concentration be after 1.0 day?

Strategy Convert the time given to seconds to match the units of the rate constant and then use the first-order integrated rate law to find [O3] at the desired time. (Note that mM is just millimolar, or 1023 M.)

Solution 24 h 60 s 60 min 1.0 day q—r q—r q—r 5 8.6 3 104 s 1 min 1 day 1h [O3] 5 [O3]0e2kt 5 5.0 mM (e2(1.0 3 10

25)(8.6

3 104)

)

5 2.1 mM

Discussion This problem reminds us of the importance of using consistent units. The argument of an exponent can never have units, so when we put in values for k and t, they must have the same time unit (seconds or minutes, for example).

Check Your Understanding In the upper stratosphere, at an altitude of about

40 km, this same reaction proceeds with a rate constant of 1.0 3 1023 s21. (The difference is caused by the higher intensity of UV radiation at higher altitude.) If an experiment were to simulate these conditions, how long would it take for an initial ozone concentration of 5.0 mM to fall to 2.1 mM?

Figure 11.7 shows a plot of [O3] versus time for the experiment described in Example Problem 11.5. A quick glance at Figures 11.6 and 11.7 shows a glaring difference between the two plots of concentration versus time: one is linear and the other is clearly not. The curve shown for the first-order reaction in Figure 11.7 is an example of what is known as exponential decay. You are probably familiar with this type of curve from

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Figure 11.7 ❚ This plot of [O3] vs. time shows that the dissociation of ozone does not follow a zero-order rate law because the graph is clearly nonlinear.

[O3]

For a first-order reaction, a plot of concentration vs. time is not linear.

Time

math classes, and you will see many other physical phenomena that show this kind of behavior as you continue your engineering studies. Suppose that we had measured the data shown in Figure 11.7, but we did not yet know the order of the reaction. The fact that the plot of concentration versus time is not linear tells us that the reaction is not zero order; it doesn’t fit the model for that rate law. Next we might ask if the reaction is first order. We can answer this question by comparing the data to the model of first-order behavior. One way to do this is to attempt to fit an exponential decay (Equation 11.6) to the data. This might be done using a graphing calculator, a spreadsheet, or some other tool that can fit nonlinear functions. If we are examining the data visually, though, a curved plot is not as easy to judge as a straight line, so we might want to manipulate the data so that a straight line results. This requires only a little algebra to arrange the integrated rate law into a linear form.

Transforming data to give a linear plot is a convenient way to compare the data with a model.

ln[A] 5 2kt 1 ln[A]0 Now we have an equation that provides a straight line when we plot ln[A] versus time. Figure 11.8 shows the same data as Figure 11.7, manipulated to give such a straightline plot. At this point, we can see that (1) if a plot of [A] versus time is linear, the reaction is zero order; and (2) if a plot of ln[A] versus time is linear, the reaction is first order. From the form of the integrated rate law equations for these cases, you should also be able to see that the slope of the linear plot must be equal to 2k. So we can find both the rate constant and the reaction order from our graph. Next we will examine the correct model for second-order kinetics.

ln[O3]

For a first-order reaction, plotting the natural log of the concentration vs. time gives a straight line. The slope of that line is equal to –k.

The quantity ln[A] has no units, so the units of the slope—and of the rate constant—are simply 1/time.

Figure 11.8 ❚ The first-order integrated rate law predicts that a plot of the natural logarithm of reactant concentration versus time should be linear. Here we see that this is the case for the dissociation of ozone.

Time


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Second-Order Integrated Rate Law For second-order reactions, the mathematical form of the integrated rate law is quite different from that we used for first-order reactions. For the consumption of reactant A, the integrated rate law is 1 [A]t

1 [A]0

— 2 — 5 kt

(11.7)

As before, [A]t is the concentration of X at time t and [A]0 is the initial concentration. If we know the value of k, we can use this equation and the same strategies used in Example Problem 11.5 to consider second-order reactions. This equation provides us with a mathematical model for a second-order reaction as a function of time. And as we’ve seen in the earlier cases, a good way to compare data to such a model is to rely on a straight-line plot. The form of the integrated rate law for second-order kinetics immediately suggests the nature of a linear plot; in this case, we must plot 1/[A] versus time. If such a plot is linear, then the reaction is second order. Example Problem 11.6 provides practice at using graphical techniques to determine the order of reactions.

E XAM PL E PRO BLEM

11.6

Among the possible fates of NO2 in atmospheric chemistry is the decomposition reaction to form NO and O2. This reaction was studied at 370ºC by a student, and the following data were obtained. Time (s)

[NO2] (mol L21)

0.0

0.3000

5.0

0.0197

10.0

0.0100

15.0

0.0070

20.0

0.0052

25.0

0.0041

30.0

0.0035

Based on these data, determine the order of this reaction and the rate constant.

Strategy The available data are the reactant concentration as a function of time for a single experiment, so we will need to use graphical techniques to determine the order of the reaction. There are three possibilities we can explore using the integrated rate laws we’ve examined. The reaction could be zero order, first order, or second order with respect to NO2. We will need to manipulate and plot the data in various ways to determine whether there is a good fit with any of these models. (Other orders are also possible, so we should be aware that all three tests could conceivably fail.) With a spreadsheet or a graphing calculator, such manipulation of data is easy. For this example, first we will calculate all of the data needed for all three plots and then make the appropriate graphs to find the linear relationship and determine the rate law.

Solution The three versions of concentration we need to plot versus time are [NO2] (zero-order model), ln[NO2] (first-order model), and 1/[NO2] (second-order model). The table on the following page gives the data we will need to plot.

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Time

[NO2]

ln[NO2]

1/[NO2]

0.0 5.0 10.0 15.0 20.0 25.0 30.0

0.3000 0.0197 0.0100 0.0070 0.0052 0.0041 0.0035

21.20 23.93 24.61 24.96 25.26 25.50 25.65

3.33 50.8 100. 140. 190. 240. 290.

First, plot concentration versus time to determine whether the reaction is zero order. 0.35

[NO2] (mol/L)

0.30 0.25 0.20 0.15 0.10 0.05 0.00

0

5

10

15 Time (s)

20

25

30

This plot is clearly not linear, so we must proceed to the other possibilities. To check for first-order kinetics, we need to plot ln[NO2] versus time.

0

ln[NO2]

–1 –2 –3 –4 –5 –6

0

5

10

15 Time (s)

20

25

30

Just like the plot for zero-order kinetics, this plot is not linear, so the reaction is not first order. Our final option is to plot 1/[NO2] versus time to see if the reaction is second order. 350

1/[NO2] (L/mol)

300 250 200 150 100 50 0

0

5

10

15 Time (s)

20

25

30


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Linear regression functions in a calculator or spreadsheet can determine the slope of a best-fit line.

This plot is linear, so the reaction must be second order. To determine the rate constant, we need to determine the slope of the line. Estimating from the graph, the line crosses 50 at 5 seconds and 200 at 21 seconds. Calculating the slope (rise/run) from those numbers, we have 150/16, which is 9.4 L mol21 s21. (The units are also given by the rise/run concept from the graph.)

Discussion Because this is an introductory text, we have chosen our examples so that they will fit one of these three simple models. In some cases, however, reaction kinetics can be quite complicated and the appropriate model might be none of the three we tried here. If none of these three plots were linear, we would have to conclude that the reaction was not zero, first, or second order. Similar integrated rate law models can be derived for other cases, but this is beyond the scope of an introductory class.

Check Your Understanding N2O5(g) can decompose to form NO2 and O2. The following data were obtained for this decomposition at 300ºC. Determine the order of the reaction and the rate constant. Time (s)

[N2O5] (mol L21)

0.0

0.1500

300.0

0.0729

600.0

0.0316

900.0

0.0203

1200.0

0.0108

1500.0

0.0051

1800.0

0.0020

Half-Life The half-life of a reactant is the time it takes for its concentration to fall to one-half its original value. Although this quantity can be defined for any reaction, it is particularly meaningful for first-order reactions. To see why, let’s return to the system we considered in Example Problem 11.5, the photodissociation of ozone by UV light in the upper atmosphere. O3 1 hn : O 1 O2 We said that this is a first-order reaction with a rate constant of 1.0 3 1025 s21 at an altitude of 10 kilometers above the earth’s surface. Suppose that we designed a laboratory experiment in which the temperature and light levels at that altitude could be reproduced. Figure 11.9 shows the type of results we might obtain. The experiment begins with PO3 5 1 atm and drops off as time goes by. The dashed lines in the figure show that the partial pressure reaches one-half its initial value after about 19 hours. The reaction continues, and the pressure falls further, reaching 0.25 atm, or one-fourth of its initial value after about 38 hours. In the first 19-hour period, the ozone pressure decreased by one-half its value from 1 to 0.5. Between about 19 hours and 38 hours on the graph, the pressure again decreased by one-half, this time falling from 0.5 to 0.25. For the particular conditions of this experiment, the ozone pressure decreases by a factor of 2 for every 19 hours. So the half-life is constant, no matter how much ozone we start with. We can obtain a mathematical expression for the half-life of a first-order reaction by substituting in the integrated rate law (Equation 11.5). By definition, when the

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1

Figure 11.9 ❚ Ozone pressure as a function of time in an experiment designed to model the destruction of ozone in the stratosphere. The ozone pressure falls by a factor of 0.5 every 19 hours.

P(O3) (atm)

0.8

0.6 0.4

0.2

0

0

20

40

60

80

100

Time (h)

reaction has been proceeding for one half-life (t1/2), the concentration of the reactant 1 must be [A] 5 —2 [A]0. Thus we have 1

0 [A]0 2 ln q—r 5 ln q— r 5 2kt

—[A]

[A]0

2[A]0

1/2

which simplifies to 1 ln — 52kt1/2 2 ln 1 2 ln 252kt1/2 or ln 25 kt1/2 Solving this equation for the half-life gives 0.693 ln 2 t1/2 5 — 5 — k k

(11.8)

Equation 11.8 relates the half-life of any first-order reaction to its rate constant. Because k does not depend on the amount of substance present, neither does t1/2. The half-life is most often used to describe the kinetics of nuclear decay. All radioactive decay processes follow first-order kinetics, and people working with nuclear processes typically use half-lives rather than rate constants to discuss these systems. According to Equation 11.8, we can easily convert the half-life to the rate constant (or vice versa), so knowing either of these values gives us the same amount of information about the reaction kinetics. (For nuclear decay, the half-life is also independent of temperature.)


Radioactive decay and its kinetics are discussed in detail in Chapter 14.

11. 7

The rate of the photodissociation of ozone in the example shown in Figure 11.9 may seem slow. But it is actually tremendously faster than what we would see in the absence of ultraviolet light. The rate constant, k, for the thermal decomposition of ozone in the dark at 25ºC is just 3 3 10226 s21. What is the half-life of ozone under these conditions?

Strategy We can see from the units of the rate constant that this is a first-order reaction. So we can use the equation for the first-order half-life.


365

Solution 0.693 0.693 t1/2 5 — 5 — 5 2 3 1025 s k 3 3 10226 s21

Analyze Your Answer A number this large may seem disconcerting. But if we stop to realize that the rate constant and the half-life are inversely related to each other, such a large number makes sense. If the rate constant is small, the half-life must be long.

Discussion Comparing this result with the half-life of 19 hours (or about 68,000 s) we saw above shows the importance of UV photons in the photodissociation reaction. In the absence of light, the rate constant is smaller by more than 20 orders of magnitude, making the half-life tremendously long.

Check Your Understanding The half-life of the photochemical decomposition of N2O5 in the stratosphere (at 40 km) is roughly 43,000 s. What is the rate constant for the reaction?

11.5

Temperature and Kinetics

The most widely recognized causes of ozone depletion in the stratosphere are chemicals that many of us still have in our refrigerators—chlorofluorocarbons (CFCs). Freon-12® is the registered trademark for CCl2CF2, the particular chlorofluorocarbon that has been most widely used in refrigerators and air conditioners. Have you ever asked yourself why you need a refrigerator? Aside from keeping your favorite beverage refreshingly cold, what advantages does the cooled storage space provide? The answer lies in the temperature dependence of the rates of chemical reactions. Without a refrigerator, food would spoil more quickly. Maintaining healthful foodstuffs is a critically important use for refrigeration. It works because at cooler temperatures, chemical reactions—including those that allow the growth of dangerous bacteria in food—are slower. You can measure this temperature effect with a simple experiment. Set up one sample of milk at room temperature and another sample in the refrigerator. Determine the time it takes the milk to smell sour. It is pretty obvious how this will turn out: the room temperature milk will sour much faster. Why does temperature affect reaction kinetics, and how can we quantify that effect?

Temperature Effects and Molecules That React We discussed the kinetic-molecular theory in Section 5.6.

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We can investigate the role of temperature in kinetics by referring to the most useful model for describing the motion of molecules—the kinetic-molecular theory. Let’s consider the events that lead to a chemical reaction between two gas-phase molecules. (The same ideas apply to reactions involving solids and liquids, too. The gas-phase case is just easier for us to discuss.) The kinetic-molecular theory says that molecules interact with one another only through collisions. So, for two molecules to react, first they must collide. We should not expect that all collisions between molecules result in chemical reactions, though. Our experience suggests that the air surrounding us is stable, even though the N2 and O2 molecules are constantly colliding. So when nitrogen and oxygen collide, the probability of reaction must be very low. If collisions alone cannot explain reactivity, what additional variables come to mind? One possibility is the kinetic energy of the colliding molecules. Slow moving molecules collide with gentle impacts that do not usually result in reaction, whereas fast moving molecules seem more likely to react.


Figure 11.10 ❚ The two curves

Number of molecules

At lower temperatures, molecular speeds are lower. Low T At higher temperatures, more molecules will be moving with the higher speeds needed to produce reactive collisions.

High T

This is the minimum velocity needed to produce a collision with sufficient kinetic energy to allow reaction.

show the Maxwell-Boltzmann distribution of molecular speeds for two temperatures. The shaded areas represent the number of molecules traveling fast enough to collide with sufficient energy to react. Note that there are many more fast moving molecules at higher temperature.

Velocity

Recall from Section 5.6 that the Boltzmann distribution describes the speeds of molecules. Figure 11.10 shows the speed distributions for gas molecules at two temperatures. At the higher temperature, a larger fraction of the molecules moves at high speeds. Thus at higher temperatures, reactions tend to proceed more quickly because more of the collisions between reactants have the high energies needed to bring about reaction. Why do reactions require high-energy collisions? Chemical reactions involve the breaking and forming of chemical bonds. From thermodynamics, we know that breaking bonds requires energy input and forming bonds liberates energy. Intuition suggests that before new bonds can form, old ones must break, or at least begin to break. So, many chemical reactions need energy input at the outset to start breaking bonds in reactants. Figure 11.11 provides one way to visualize this requirement. Standard thermodynamic data allow us to predict whether a reaction is exothermic or endothermic, as we saw in Chapter 9. But to get from reactants to products, the reaction must first overcome an energy threshold, called the activation energy or activation barrier. Collisions of fastmoving particles provide sufficient kinetic energy to overcome the activation barrier. The larger this barrier, the more kinetic energy will be needed. So the activation energy is what dictates how energetic a collision must be to lead to reaction. Collision geometry also plays a role in whether a collision between fast-moving molecules is effective. As you can see in Figure 11.12, a collision between a molecule of N2O and an oxygen atom may or may not lead to a reaction, depending on the

Activation energy

The activation energy is sometimes called an activation barrier because on a graph like this it looks like a barrier that the reactants must overcome to produce products.

plot for an exothermic chemical reaction. To progress from reactants to products, the molecules must collide with enough energy to pass over the activation barrier.

Energy

Reactants

Figure 11.11 ❚ Potential energy

Products

Reaction progress is a measure of the extent of bond breaking and forming as a reaction proceeds.

Reaction progress


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If an O atom strikes the nitrogen end of an N2O molecule, it cannot react to form molecular oxygen.

‡

If the O atom strikes the oxygen end of an N2O molecule, it is able to form an O O bond, leading to N2 and O2 as products.

Figure 11.12 ❚ Geometric factors can also be important in determining whether a molecular collision is reactive. For the oxygen atom to react with N2O, it must strike the oxygen end of the molecule. The importance of collision geometry depends on the shapes of the molecules in the reaction.

More specifically, the activated complex is the highest energy point along the minimum energy path to products.

geometry of the collision. At the moment of an effective collision, both bond breaking and bond formation are occurring. For an instant, as bond rearrangement is occurring, an unstable intermediate species, called an activated complex, exists in the reaction mixture. (This activated complex is denoted ‡ in Figure 11.12.) The activated complex represents the highest energy point along the route from reactants to products. Because it is highly unstable, the activated complex may have a lifetime as short as 10215 s. Formation of an activated complex and its subsequent decomposition into reaction products depends on both the energy and the geometry of the collision.

Arrhenius Behavior Svante Arrhenius also studied acids and bases. The notion that acids form H1 and bases form OH2 in water is attributed to him.

The temperature dependence of chemical kinetics has been studied for many years. For most reactions, a relationship known as the Arrhenius equation can be used to describe the temperature dependence of the rate constant, k: k 5 Ae2Ea/RT

(11.9)

Here Ea is the activation energy, R is the universal gas constant, T is the temperature (in kelvins), and A is a proportionality constant called the frequency factor or preexponential factor. We noted earlier that a large activation energy should hinder a reaction. Equation 11.9 shows this effect: As Ea increases, k will be smaller, and smaller rate constants correspond to slower reactions. Note that because the temperature and activation energy appear in the exponent, the rate constant will be very sensitive to these parameters. That’s why fairly small changes in temperature can have drastic effects on the rate of a reaction. The Arrhenius equation can be used to determine activation energy experimentally. Temperature is a parameter that we can usually control in an experiment, so

368


Intercept = ln A

–Ea R

ln k

Slope =

1/T (K–1)

Figure 11.13 ❚ According to the Arrhenius equation, a plot of ln k vs. 1/T should give a straight line with a slope equal to 2Ea/R. Such a plot is frequently used to determine the activation energy of a reaction.

it may help to remove it from the exponent. We can do this by taking the natural logarithm of both sides of Equation 11.9: 2Ea ln k 5 — 1 ln A RT Rearranging, we get 2Ea 1 ln k 5 — q—r 1 ln A R T This shows us how we can determine values for Ea or A. We will need to measure k for at least two temperatures, but we already know ways to do that. The rearranged equation is in the familiar form of a line, y 5 mx 1 b, so a plot of ln k versus 1/T should be a straight line. The line’s slope is 2Ea/R, where R is the gas constant, whose value is known. So we can get an estimate of the activation energy from the slope, as shown in Figure 11.13. This type of problem is easily attacked with a graphing calculator or a spreadsheet, and there are several such problems at the end of this chapter.


When working with the Arrhenius equation, the most commonly used form for R is 8.314 J K 21 mol21. It is essential that all units are chosen so that the quantity (Ea /RT) is unitless.

11. 8

In the troposphere, ozone can be converted to O2 by the following reaction with hydrogen oxide radicals: HO ? (g) 1 O3(g) : HO2?(g) 1 O2(g) The following values for the rate constant, k, for this reaction were measured in experiments at various temperatures. k (L mol21 s21)

Temperature (K)

1.0 3 107

220

7

5.1 3 10

340

1.1 3 108

450


369

(a) Does this reaction display Arrhenius behavior? (b) Estimate the activation energy from these data.

Strategy To determine whether or not the reaction displays Arrhenius behavior, we will need to compare the data to the model represented by the Arrhenius equation. To do this we plot ln k versus 1/T. If the plotted data form a straight line, the slope of the line is 2Ea/R. So, we can determine the activation energy from that slope.

Solution First, expand the table to include the data we need to plot. k (L mol21 s21)

ln k

Temperature (K)

1/T (K21)

1.0 3 107

16.1

220

4.5 3 1023

5.1 3 107

17.7

340

2.9 3 1023

1.1 3 108

18.5

450

2.2 3 1023

Now plot ln k ( y axis) versus 1/T (x axis) The quantity ln k is considered unitless since the logarithm of a unit has no physical meaning. So the units of the slope of this graph will be kelvins.

19.0 18.5

ln k

18.0 17.5 17.0 16.5 16.0 15.5

0

0.001

0.002 0.003 1/T (K–1)

0.004

0.005

The data do appear to fall on a straight line. So our answer to question (a) is that yes, the reaction does exhibit Arrhenius behavior. (It would be better if we had more than three data points on which to base this conclusion.) To determine the slope of the straight line, we could use the curve-fitting routine built into a graphing calculator or spreadsheet program. Or we can determine the slope from two points on the line. If our straight line for these data includes the first and last points, which the graph suggests as a reasonable estimate, we can use data from our expanded table directly. 18.5216.1 Rise Dln k Slope 5 — 5 — 5 —— 23 Run 1 (2.2 3 10 2 4.5 3 1023) K21 D q— r T Ea 5 21.04 3 103 K 52 — R We can insert

R 5 8.314 J K21 mol21 and solve for Ea. Ea 5 (1.04 3 103 K) (8.314 J K21 mol21) 5 8.7 3 103 J mol21 5 8.7 kJ mol21

Analyze Your Answer To assess this answer, we need some familiarity with the magnitude of activation energies. Typical Ea values are on the order of tens of kJ mol21, so this is slightly small. But free radicals are very reactive, and that is consistent with a small activation energy.

370


Discussion Often the best-fit line for experimental data does not explicitly cross through any of the individual data points. Graphing calculators or spreadsheet programs provide the best and simplest way to determine the slope of such a line. But it is important to keep in mind that the slope calculation is only meaningful if the straightline model provides a reasonable fit to the data. It is always a good idea to look at the data and the best-fit line on the same graph to verify that you are satisfied with the quality of the fit. If the data do not suggest a linear relationship, then we should not put our confidence in the meaning of a slope or other parameter taken from a linear fit.

Check Your Understanding Once generated, HO2· may react with carbon monoxide to reform the hydrogen oxide radical and carbon dioxide. This reaction plays a role in ozone pollution in major cities. Assume that an experiment to determine the rate constant of this reaction provides the data tabulated here. What is the activation energy of this reaction? k (L mol21 s21)

Temperature (K)

5.9

301

4.7 3 10

4

650

For most students, the most challenging mental connections that are forged in general chemistry are those between bulk macroscopic properties and molecular level origins of those bulk properties. The Arrhenius equation provides us with a good opportunity to focus on this type of mental connection. k 5 Ae2Ea/RT

(11.9)

There are two important factors in this equation—the argument of the exponent and the frequency factor, A. Figure 11.14 shows how these two components of the equation arise from considerations at the molecular level. The figure shows a reaction in which two molecules of ClO are converted to Cl2 and O2: 2 ClO : Cl2 1 O2 In keeping with our usual color code, the figure shows chlorine atoms in green and oxygen atoms in red. This particular reaction was chosen because it is related to the atmospheric chemistry of ozone. But the same arguments can be applied to any other reaction, too. Perhaps the effect of temperature on the Boltzmann distribution is familiar by now, but we show it for two separate cases—low temperature (2) and higher temperature (3). We have noted previously that Boltzmann distributions imply that more molecules can react at higher temperatures. In the final three illustrations, we depict the connection between the molecular velocities that are incorporated in (2) and (3) and reactivity. Drawing (4) shows that the slow-moving molecules have insufficient energy to react; the collision energy does not exceed the activation energy. In (5), however, the high velocity molecules have large collision energies—in excess of the activation energy—and reaction occurs. We could just as easily have drawn this effective collision originating from the higher temperature Boltzmann distribution, where even more molecules will undergo these high-energy collisions. The accompanying MathConnections section examines how the equation makes these ideas quantitative.


371

1 A reacting sample of CIO contains many molecules that collide with one another quite often. This illustration represents an incredibly tiny fraction of the molecules that would be present in the macroscopic sample.

At higher temperature

2 At low T, the distribution of molecular velocities tends toward lower values.

Fraction of molecules

Fraction of molecules

At lower temperature

3 At higher T, the distribution of molecular velocities tends toward higher values.

Velocity

Velocity

Activation energy

Activation energy

Activation energy

Collision energy

Collision energy

Collision energy

4 Slower moving molecules do not have enough collision energy to react.

5 With sufficient collision energy, reactions may occur.

6 Some collisions, like the one depicted here, have enough energy but do not lead to a reaction. The preexponential factor of the Arrhenius equation helps to factor this type of event into the average behavior of molecules that results in observed bulk properties.

Figure 11.14 ❚ The connection between temperature, collision energy, and reaction rate is developed here. The activation energy establishes a minimum threshold that must be overcome before a collision can lead to reaction. The number of energetic molecules increases with temperature, according to the Boltzmann distribution.

372


MathConnections Figure 11.14 shows how important it is for the collision energy to exceed the activation barrier, but how is this concept expressed in the Arrhenius equation? The argument of the exponential function is the ratio of these two values. Recall from Section 5.6 where we introduced gas kinetic theory that the average velocity (the root-mean-square velocity, to be more precise) is given by the equation cvrms 5

√ MM d ‾‾‾ 3RT

—

MM is the molar mass of the gas. If we square both sides of this equation, we can isolate RT: cRT 5 — MMv 2rmsd

1 3

Noting that the molar mass is the mass of a molecule times Avogadro’s number (NA), we can see that NA RT 5 — mv 2rms 3 1 Because kinetic energy is expressed as — mv2, RT is directly related to the average 2 kinetic energy: 2NA 2NA 1 RT 5 — q— mv 2rmsr 5 — KEavg 3 2 3 Thus, the macroscopic Arrhenius equation is based on the average kinetic energy of the molecules in a sample. The Boltzmann distribution can be evaluated to quantify that average behavior, and the connection between macroscopic and microscopic is made. The Arrhenius equation contains a ratio of Activation energy

—

Kinetic energy

When we change the kinetic energy part of this ratio by changing the temperature, the activation energy remains the same. The frequency factor, A, accounts for the collision frequency, which will depend on the sizes of the molecules involved. It also reflects the fact that even at high temperatures and high collision energies, not all collisions result in reaction, due to geometric effects like those illustrated in Figure 11.12. For the reaction considered in Figure 11.14, an effective collision might require that the oxygen ends of the two ClO molecules collide directly with one another, for example, to allow the O"O bond to begin to form.

11.6

Reaction Mechanisms

Earlier, we noted the net reaction of the Chapman cycle: 3 O2 : 3 O2. In some ways, it appears pointless to talk about the kinetics of a reaction that has no net change. But as we have already seen, the importance of the Chapman cycle lies in the details. Without the Chapman cycle, the upper atmosphere would not have an ozone layer, and the quantity of harmful UV radiation reaching the surface of the planet would be significantly greater. This natural cycle points to the importance of the mechanism of the reaction. A reaction mechanism is a collection of one or more molecular steps that account for the way reactants become products. If the overall equation is like a parts list and a finished product, then the mechanism is like the assembly instructions.

The study of reaction mechanisms on a molecular level is referred to as chemical dynamics.


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Table

❚

11.1

Summary of molecularity of elementary reactions

Type of Elementary Reaction

Molecularity

Rate Law

A : products

Unimolecular

Rate 5 k[A]

A 1 B : products

Bimolecular

2 A : products A 1 B 1 C : products

Termolecular

2 A 1 B : products A 1 B 1 C 1 D : products

Rate 5 k[A][B] Rate 5 k[A]2 Rate 5 k[A][B][C] Rate 5 k[A]2[B]

Not observed

In many cases, such as the Chapman cycle, if we want to understand the kinetics of the process we must consider the reaction mechanism. What is the stepwise process of bond breaking and bond formation behind the overall reaction?

Elementary Steps and Reaction Mechanisms

It is very important to remember that we can use the stoichiometric coefficients as the reaction orders only when we are considering an elementary step!

We have already reasoned that chemical reactions occur when reactants collide. Based on the kinetic-molecular theory, collisions between two particles should be rather common. Three-particle collisions are much less likely because they require that three randomly moving species arrive at the same place at the same time. And four-particle collisions are virtually unheard of. However, when we look at the stoichiometric coefficients of chemical equations, they seem to imply that collisions can occur among many particles: balanced equations often involve stoichiometric coefficients far greater than two or three. To understand reaction mechanisms, we need to distinguish between the overall stoichiometry of a reaction and the steps in the reaction mechanism. The individual steps in a reaction mechanism are called elementary steps. Unlike the overall stoichiometric equation, the coefficients for reactants in an elementary step do provide the exponents in the rate law for that step. According to our reasoning above, only three types of elementary steps are likely to occur, those involving one, two, or three molecules. Steps with one reactant are called unimolecular, and those with two and three reactants are called bimolecular and termolecular, respectively. The molecularity tells us the overall order of the rate law for the elementary step, as summarized in Table 11.1. Let’s consider a particular reaction mechanism so that we can define some other terms. The destruction of ozone by chlorine radicals is expressed by the net reaction, 2 O3 : 3 O2. The accepted mechanism for this reaction is Cl? 1 O3 : ClO? 1 O2 ClO? 1 O3 : Cl? 1 2 O2 Net:

2 O3 : 3 O2

Notice that the equations for elementary steps look just like those for any other chemical reaction. So it is important to remember that when we write a reaction mechanism, each equation is always an elementary step. In the mechanism above, you can see two important characteristics of reaction mechanisms: An intermediate in the overall reaction will be a product in one elementary step and a reactant in another.

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1. The ClO generated in the first reaction is consumed in the second reaction. A chemical species that is generated in one step and consumed in a later step is called an intermediate or a reactive intermediate. Many mechanisms involve one or more intermediates.


2. When the steps of the mechanism are properly summed, they give the observed stoichiometry of the overall reaction. Because of the enormous number of collisions occurring at any moment in a gas-phase reaction and the rapid formation and decomposition of reactive intermediates, reaction mechanisms can never be known with certainty. However, by observing reaction rates and comparing proposed mechanisms with experimental rates and with the overall stoichiometry, we can propose a mechanism that fits available data. If more than one mechanism seems plausible, we may be able to design additional experiments to try choosing between them. Example Problem 11.9 provides some practice in identifying valid mechanisms and the terms associated with them.


11. 9

The decomposition of N2O5 is given by the equation 2 N2O5(g) : 4 NO2(g) 1 O2(g) The following mechanism is proposed for this reaction: N2O5 : NO2 1 NO3 NO2 1 NO3 : NO2 1 NO 1 O2 NO 1 NO3 : 2 NO2 (a) Does this mechanism as written provide the correct stoichiometry? If not how does it need adjustment? (b) Identify all intermediates in the mechanism. (c) Identify the molecularity of each step in the mechanism.

Strategy Consider the definitions for reaction mechanism and see how this example of a mechanism illustrates those definitions.

Solution (a) A simple sum of the mechanism yields N2O5 1 NO3 : O2 1 3 NO2 This is not the desired stoichiometry, but if we multiply the first step by 2, we have 2 N2O5 : 2 NO2 1 2 NO3 Adding that to steps 2 and 3, 2 N2O5 : 2 NO2 1 2 NO3 NO2 1 NO3 : NO2 1 NO 1 O2 NO 1 NO3 : 2 NO2 yields 2 N2O5 : O2 1 4 NO2 So the mechanism does provide the correct stoichiometry, as long as we realize that some steps may need to occur more often than others. (b) The intermediates in this case are NO3 and NO. (c) The first step is unimolecular, whereas steps two and three are bimolecular.

Discussion This type of problem solving is really performed in reverse. We knew what the target reaction had to be (the end of the problem) and had to go back to other data given to solve the problem. Reaction mechanisms certainly aren’t the only


375

time we use this problem solving strategy in general chemistry. Stop for a moment and try to think of another type of chemistry problem where working the problem in reverse makes the most sense.

Check Your Understanding The following mechanism was once proposed for the decomposition of N2O5: N2O5 : N2O3 1 O2 N2O3 : NO2 1 NO NO 1 N2O5 : 3 NO2 (a) Is this mechanism plausible? (b) Identify the intermediates in the reaction. (c) Determine the molecularity of each step.

Mechanisms and Rate: The Rate-Determining Step For an elementary step, the orders of reaction are given by the coefficients in the balanced equation.

Once we have proposed a mechanism, it is easy for us to write the rate law for each elementary step in that mechanism. What we can measure experimentally, however, is the rate law for the overall reaction. How can we relate the observed overall rate law to the rate laws for the individual elementary steps in a mechanism? The answer lies in the fact that most mechanisms include one step that is much slower than all the others. This slow step is called the rate-determining step and it dictates the rate law. Why does the slowest step assume this role? Consider this analogy. Suppose you eat lunch at a very popular cafeteria where the food is great and the prices are low. The only problem is that the service is a little slow. As the line moves along, it takes each person 5 minutes to get food and then another 30 seconds to pay. If we watch the cash register in this cafeteria, we can expect to see one customer walking away with food every 5 minutes—the rate of the slow step in the “mechanism” for getting lunch. We can take this analogy one step further. Suppose that the owner is concerned that the line is too long and is driving customers away. To speed things up, he buys a new cash register that cuts the paying time down to just 15 seconds per customer. Much to his surprise, he finds that one customer still walks away from the register every 5 minutes! Speeding up the fast part of the process has no impact on the overall rate because that rate is set by the slow step. That’s why we call the slow step in a reaction mechanism “rate determining.” In a chemical reaction, the fast steps may be several orders of magnitude faster than the slow step. In the formation of ozone through the Chapman cycle, the slow step is the last one: the reaction of ozone with atomic oxygen. O 1 O3 : 2 O2 It is because this step is rate determining for the cyclic reaction that there is a protective ozone layer in the stratosphere. Once oxygen atoms combine to form ozone, the O3 survives long enough to absorb UV radiation, and that absorption prevents the UV light from reaching the surface of the planet. If the last step were not slow, the ozone created earlier in the cycle would be destroyed chemically before it encountered a photon of UV light. Without the ozone to absorb it, that light would pass through the stratosphere to harm living organisms in the troposphere.

11.7

Catalysis

We noted from the outset of this chapter that environmental concerns about ozone depletion arise from the increase in the rate of destruction of ozone in the stratosphere. This is a process apparently caused by man-made chemicals. Specifically, chlorofluorocarbons used as refrigerants have been cited as a cause of ozone depletion. Yet a look at the Chapman cycle does not show any obvious role for CFCs. So how can these 376


molecules accelerate the destruction of ozone if they do not even appear in the underlying chemical equations? Catalysis is a process in which a reaction rate is influenced by the presence of substances that are neither reactants nor products in the overall equation. A catalyst is a substance that increases the rate of the reaction but is neither created nor destroyed in the process. How can CFCs catalyze ozone depletion?

Homogeneous and Heterogeneous Catalysts Catalysts can be divided into two broad categories. Homogeneous catalysts are those that are in the same phase as the reacting substances, whereas heterogeneous catalysts are in a different phase from the reacting species. For gas-phase reactions, the heterogeneous catalyst typically is a solid surface. Both types of catalysis are important in atmospheric processes. The catalytic destruction of ozone in the stratosphere involves reactions between gases there, so it is an example of homogeneous catalysis. The most important catalyst for this process is chlorine. Much of the chlorine present in the stratosphere comes from CFC molecules that were released in the troposphere and slowly migrated to the stratosphere. (Because they are very unreactive at ground level, nearly all CFCs that are released into the atmosphere eventually find their way to the stratosphere.) Upon absorption of UV light, the CFCs initiate a catalyzed reaction mechanism: Step 1:

CF2Cl2 1 hn : CF2Cl 1 Cl?

Step 2:

Cl? 1 O3 : ClO? 1 O2

Step 3:

ClO? 1 O3 : Cl? 1 2 O2

We encountered the second and third reactions earlier in our discussion of mechanisms. The chlorine atoms serve as a catalyst for ozone decomposition because (a) they are not part of the reaction stoichiometry, (b) they are not consumed by the reaction, and (c) they increase the rate of the net reaction. Because the chlorine atom is regenerated in Step 3, it can catalyze another reaction. An individual chlorine atom may cycle through this mechanism often enough to destroy 100,000 or more ozone molecules. Because of this, the U.S. banned the use of CFCs as propellants in aerosol cans in 1978. The Montreal Protocol, an international treaty, calls for an end to CFC production. Non-CFC refrigerants and propellants are now in use in new equipment, although old refrigerators and air conditioners still contain CFCs that may one day be released to the environment. Heterogeneous catalysis has a role in the atmospheric chemistry of ozone in the troposphere as well. Catalytic converters in automobiles are filled with a porous ceramic material, which provides a surface that catalyzes the removal of CO and NOx (nitrogen oxides) from the exhaust. (Nitrogen oxides initiate the formation of ozone and other lung irritants in photochemical smog. We will examine this process in detail in Section 11.8.) The process by which catalytic converters operate is shown in Figure 11.15. The types of steps shown there are found in most examples of heterogeneous catalysis.

Dr. Susan Solomon’s work at the National Oceanic and Atmospheric Administration shows that these reactions occur on the surface of ice particles.

1. The reactive species adsorb, or stick, onto the surface of the catalyst. (When a molecule is adsorbed, it sticks to the surface. When a molecule or substance is absorbed, it penetrates into the bulk of a material, like water into a sponge. Desorption is the opposite of adsorption.) 2. Species migrate on the surface until they encounter each other. 3. The reaction takes place on the surface. 4. The products desorb from the catalyst surface.

Molecular Perspective of Catalysis We noted earlier that the activation energy of a chemical reaction plays a critical role in its kinetics. Catalysts increase the rate of reaction by providing a new reaction 11.7

Catalysis

377

2 …adsorb onto the platinum surface…

1 NO molecules collide with the catalyst…

4 Pairs of adsorbed N and O atoms react to form N2 and O2.

3 …and dissociate. N and O atoms bond with Pt atoms.

5 Products desorb, and the surface can catalyze further reaction.

Figure 11.15 ❚ A schematic model of the steps involved in the catalytic removal of NO from exhaust gas.

pathway that lowers the activation energy. Figure 11.16 illustrates this concept for the conversion of two ozone molecules to three oxygen molecules. Despite the fact that the addition of chlorine atoms requires an extra step relative to direct collisions of two ozone molecules, the net energy required for the reaction is lower. We saw in Section 11.5 that the energy needed to overcome the activation barrier is normally acquired from collisions of fast moving particles. When a catalyst is introduced, the amount of energy needed decreases, and a larger fraction of the molecules have speeds that are high enough to provide the needed energy. So the direct effect of a catalyst at the molecular level is to lower the activation energy, and the net result is an increased rate of reaction. Figure 11.16 ❚ Catalysts increase the rate of reaction by providing an alternative reaction pathway with lower activation energy. Energy

The uncatalyzed reaction must overcome a relatively high activation barrier.

2 O3 Catalysis leads to a new reaction pathway that has a lower activation energy.

3 O2 Reaction progress

378


Catalysis and Process Engineering One of the most prominent challenges of modern chemistry and materials science is the rational design of catalysts for industrial processes. Catalysts are important in industries such as petroleum refining, where billions of dollars are at stake, so the incentives for research in these areas are great. As in any design process, we can begin by listing some desirable features for a catalyst. It should have sufficient longevity or durability to last for many reaction cycles. Although the definition of a catalyst might seem to imply that the catalyst should last indefinitely, usually this is not true in practice. Many catalysts suffer from chemical or mechanical degradation and eventually become useless. A good catalyst should also have a high “turnover number.” This means that it can generate large amounts of products in short times. (Turnover number is defined as the number of molecules that can react per catalyst binding site per unit time.) Another crucial property of catalysts is selectivity. We often want the catalyst to speed up one, and only one, reaction to minimize the formation of unwanted byproducts. A highly selective catalyst will accomplish this goal. Historically, catalysts have often been chosen by trial and error. But as our understanding of reaction mechanisms, molecular structure, and materials properties has increased, many researchers feel it is now becoming possible to apply standard engineering design principles to improve catalysts.

A high turnover number is advantageous for an industrial process because it generates more product per unit time.

INSIGHT INTO

11.8

Tropospheric Ozone

Ozone in the stratosphere helps make life possible by blocking large amounts of harmful ultraviolet radiation. In the troposphere, however, it is a major lung irritant. It causes localized problems when it is formed from O2 by an electrical discharge (such as lightning or the electrostatic mechanism in a photocopy machine) or by ultraviolet light. As the major constituent of photochemical smog, ozone contributes to health problems in urban areas. Monitoring by the U.S. Environmental Protection Agency suggests that O3 is quite prominent in the troposphere. Typically, almost 2,000 violations of air quality standards are reported during summer months in North America. The kinetics of ozone formation at ground level differs substantially from that in the stratosphere. The formation of ozone in smog requires several steps. The key compound, NO2, forms in the high temperature environment of automobile engines. When NO2 absorbs sunlight, it can dissociate: NO2 1 hn : NO 1 O

We first discussed photochemical smog in Section 5.1; ideas from this chapter allow us to provide a deeper explanation.

(11.10.a)

The involvement of light in this initiating reaction provides the term photochemical smog. It is plausible to expect that the oxygen atoms generated by this reaction can then form ozone by the reaction O 1 O2 1 M : O 3 1 M

(11.10.b)

Here we include an unidentified third species, M. Such a species must be present to remove excess energy in the ozone molecule once it forms, or else that excess energy will cause the newly formed molecule to decompose immediately. Even if reaction 11.10.b occurs, the ozone that forms could be consumed quickly by reaction with NO. O3 1 NO : NO2 1 O2

(11.10.c)

Just as in the stratosphere, the balance between the rates at which ozone is formed and destroyed will determine the ozone level. The rate constants in Table 11.2 show that reaction 11.10.c is much faster than reaction 11.10.b. Thus, if no other chemicals present react with NO, ozone is not likely to form in the troposphere. The chemicals that most often give rise to the consumption of NO are volatile organic compounds (VOCs). Common VOCs are industrial solvents such as pentane,

11.8 Tropospheric Ozone

379

Table

❚

11.2

Rate constants for some reactions in atmospheric chemistry (Data adapted from NBS technical Note 866, “Chemical Kinetic and Photochemical Data for Modeling Atmospheric Chemistry,” U.S. Department of Commerce, National Bureau of Standards, 1975.)

Product ingredient lists sometimes use the general term “petroleum distillates” rather than actual chemical names such as hexane.

Reaction (at 300 K unless indicated)

Rate Constant

O 1 O2 1 M : O3 1 M

5.0 3 105 L2 mol22 s21

O3 1 NO : NO2 1 O2

1.0 3 107 L mol21 s21

O3 1 NO2 : NO3 1 O2

3.0 3 104 L mol21 s21

O3 1 OH : OOH 1 O2

4.8 3 107 L mol21 s21

O3 1 Cl : OCl 1 O2

1.1 3 1010 L mol21 s21

O 1 N2O : N2 1 O2 (1200 K)

8.1 3 105 L mol21 s21

NO 1 CH3O2 : CH3O 1 NO2

3.8 3 108 L mol21 s21

ClO 1 ClO : Cl2 1 O2

1.4 3 107 L mol21 s21

hexane, and benzene (Table 11.3). As a group, these are known as petroleum distillates because they are obtained by distilling crude oil. We can designate this class of chemicals as RH, where R can be any organic radical, such as !CH3, !C2H5, etc. The role of these RH species in smog formation begins when they react with HO? radicals. RH 1 HO? : R? 1 H2O

(11.10.d)

We have designated radical species by a dot representing an unpaired electron. The product radical, R?, can react with oxygen in a fast reaction: R? 1 O2 1 M : RO2? 1 M

(11.10.e)

The species RO2? is called an alkylperoxy radical. It removes NO and thereby increases the lifetime of any ozone that is formed. The reaction produces NO2 and another radical, RO?. RO2? 1 NO : RO? 1 NO2

Modern gas caps and pump nozzles are designed to minimize the emission of VOCs and reduce smog formation.

(11.10.f )

Representative rate constants for these reactions are included in Table 11.2. Note the wide range of rate constant values. We can see that the reactions to remove NO are roughly 40 times faster than the reaction of NO with ozone. The result is that if NO2 (the initiator) and VOCs are both present, then with adequate sunlight, smog may form and ozone can form within that smog. Table

❚

11.3

Examples of volatile organic chemicals and their sources

380

VOC

Source

Petroleum distillates (pentane, hexane, benzene)

Gasoline spills and evaporation

Terpenes

Emitted from live plants (e.g., odors of trees)

Alcohols and aldehydes

Solvents (for paint thinners, etc)

Plasticizers

Outgassing from new rugs, upholstery, electronic equipment


These few reactions do not completely describe the chemical processes that form smog. Additional reactions not shown here produce other substances that are strong lung irritants. But even with the few reactions we have examined here, we can begin to appreciate how complicated the study of atmospheric chemistry is. Many pollutants are present at concentrations of parts per million or less, and their concentrations may vary with the seasons, the weather, or even the time of day. Yet they result in real threats to human health. How can atmospheric chemists hope to unravel such a complex problem? The usual approach is to construct a model that incorporates as many of the relevant reactions as possible. Scientists then use computers to solve simultaneous rate law equations for the connected processes. One commonly used model, known as the “urban airshed model,” incorporates 86 chemical reactions among 36 different species. The role of individual reactions can be explored by adjusting their rate constants or even by removing them from the model entirely. Several of the species involved are those we have discussed in this chapter, and the major pollutant that results is ground-level ozone. Among all of the reactions included in the mechanism, ozone is produced in only one of them, and yet it is the key product of the entire mechanism. Of course, uncertainty in the values of any of the rate constants used in this model can shift the expected concentrations of many substances, sometimes dramatically. Our brief introduction to chemical kinetics cannot provide us with the tools needed to investigate such a complicated process any further. But it does provide some insight into the roles of both chemistry and mathematical modeling in understanding and confronting an important societal problem.

FO C U S ON


Question At room temperature, it takes about 3 days for a container of milk to turn sour. An identical container stored in a refrigerator turns sour only after 11 days. How would you use this information to estimate the activation energy of the chemical reaction that turns milk sour? What information would have to be looked up?

Strategy This question involves rates and temperatures, so it is a variation on the kind of kinetics experiments we addressed with the Arrhenius equation. In this case, we have to identify a few items before we can apply the equation, but the key issues are similar to example problems we have already seen. The time it takes for the milk to sour is related to the inverse of a rate constant, and the relationship between rate constant, temperature, and activation energy is found in the Arrhenius equation. 2Ea 1 ln k 5 — q— r 1 ln A R T

Solution The key to this type of question is to identify the temperature conditions with some precision. Room temperature is typically taken to be in the range of 20–25°C. If we accept this range for our temperature, we would have to report a range of possible activation energies for this reaction. The temperature of the refrigerator provides a similar problem. A typical refrigerator maintains an internal temperature of about 3–7°C. Once we establish these temperature ranges and determine a rate for the souring reaction, we can plot the rate constant versus 1/T and use the slope of this line to estimate the activation energy.

S U M M ARY Because chemical reactions are examples of change, it is only natural that understanding the rate of the change can be important. The study of chemical kinetics includes measurements of rates of reactions, as well as studies of the detailed mechanisms by which they occur.

The rate of a chemical reaction is expressed in terms of the changes in concentration per unit time, and a careful definition must account for the stoichiometry of the reaction. Reaction rates change as a reaction proceeds, so we also need to distinguish between average and instantaneous rates.

Summary

381

A number of factors influence the rate of a chemical reaction. The concentration of reactants present is one critical factor, and the experimental determination of this effect results in an expression referred to as the rate law. We can work with rate laws in either of two forms: the differential rate law and the integrated rate law. The differential rate law tells us the rate if we know the appropriate concentrations, whereas the integrated rate law predicts the concentration of a reactant as a function of time. Temperature also affects the rate of chemical reactions. An increase in temperature generally results in an increased rate of reaction. This can be understood by noting that most reactions, viewed from the molecular perspective, require some energy input to get them started. Increasing the temperature creates more

energetic molecular collisions, and the rate increases. Experimental studies of reaction rates as a function of temperature provide the data needed to measure the activation energy. For many reactions, the molecular level understanding is more complicated than asking if the reactants collide with enough energy to overcome the activation barrier. Many reactions proceed through reaction mechanisms that may comprise several elementary steps. Mechanisms are difficult to prove experimentally, but any mechanism must be consistent with experimental evidence for it to be useful as a model that explains how a reaction proceeds. Understanding reaction mechanisms helps us explain catalysis. A catalyst provides an alternative mechanism for a reaction. It lowers the activation energy and increases the rate.

KEY TERMS activated complex (11.5)

frequency factor (11.5)

rate law (11.3)

activation energy (11.5)

half-life (11.4)

rate-determining step (11.6)

Arrhenius equation (11.5)

heterogeneous catalyst (11.7)

reactive intermediate (11.6)

average rate (11.2)

homogeneous catalyst (11.7)

reaction mechanism (11.6)

bimolecular (11.6)

instantaneous rate (11.2)

reaction rate (11.2)

catalyst (11.7)

integrated rate law (11.4)

second-order reaction (11.4)

Chapman cycle (11.1)

order of reaction (11.3)

termolecular (11.6)

differential rate law (11.3)

photochemical smog (11.8)

turnover number (11.7)

elementary steps (11.6)

preexponentical factor (11.5)

unimolecular (11.6)

first-order reaction (11.4)

rate constant (11.3)

zero-order reaction (11.4)



INSIGHT INTO Ozone Depletion 11.1 What is an allotrope? 11.2 Describe what would happen to the ozone molecules if a small amount of ozone were generated and released in a room.

(b) Drying dishes by hand (c) The beating wings of a mosquito (d) Eyes blinking 11.11 For each of the following, suggest an appropriate rate unit. (a) Heart beating

11.3 In what region of the atmosphere is ozone considered a pollutant? In what region is it considered beneficial?

(b) Videotape rewinding

11.4 What are the steps in the Chapman cycle? Explain the cycle’s importance.

(d) Gas evolving in a very fast chemical reaction

11.5 What is the net chemical reaction associated with the Chapman cycle? 11.6 At what points in the Chapman cycle do photochemical reactions take place? 11.7 Which photochemical reaction in the Chapman cycle has the highest energy input via light? 11.8 Describe what is meant by the term ozone hole. Would you say that the average nonscientist pictures this issue correctly based on the phrase? 11.9 Is the ozone hole permanent? Rates of Chemical Reactions 11.10 For each of the following, suggest appropriate rate units. (a) Driving from one place to another

382


(c) Automobile wheels rotating 11.12 Rank the following in order of increasing reaction rate. (a) Dynamite exploding (b) Iron rusting (c) Paper burning 11.13 Distinguish between instantaneous rate and average rate. In each of the following situations, is the rate measured the instantaneous rate or the average rate? (a) In a hot dog eating contest, it took the winner only 4 minutes to eat 20 hot dogs, so he ate 5 hot dogs per minute. (b) At Minute 1.0, the winner was eating 6 hot dogs per minute, but at Minute 3.0, he was down to only 3 hot dogs per minute. 11.14 Candle wax is a mixture of hydrocarbons. In the reaction of oxygen with candle wax in Figure 11.4, the rate of

■


[B] for each 10-s interval from 0.0 to 40.0 s. Does the rate of change decrease from one time interval to the next? Suggest a reason for this result.

consumption of oxygen decreased with time after the vessel was covered, and eventually the flame went out. From the perspective of the kinetic-molecular theory, describe what is happening in the vessel. 11.15 In the description of the candle in Figure 11.4, we mentioned the consumption of oxygen. Assuming that candle wax is a mixture of hydrocarbons with the general formula CnH2n+2, what other variables could be measured besides the concentration of oxygen to determine the rate of the reaction? 11.16 The reaction for the Haber process, the industrial production of ammonia, is N2(g) 1 3 H2(g) : 2 NH3(g)

(b) How is the rate of change of [A] related to the rate of change of [B] in each time interval? Calculate the rate of change of [A] for the time interval from 10.0 to 20.0 s. (c) What is the instantaneous rate, D[B]/Dt, when [B] 5 0.750 mol/L? 11.21 Azomethane, CH3NNCH3, is not a stable compound, and once generated, it decomposes. The rate of decomposition was measured by monitoring the partial pressure of azomethane, in torr.

Assume that under certain laboratory conditions ammonia is produced at the rate of 6.29 3 1025 mol L21 s21. At what rate is nitrogen consumed? At what rate is hydrogen consumed? 11.17

Time (min) Pressure (torr)

Ammonia can react with oxygen to produce nitric oxide and water.

■

4 NH3(g) 1 5 O2(g) : 4 NO(g) 1 6 H2O(g) If the rate at which ammonia is consumed in a laboratory experiment is 4.23 3 1024 mol L21 s21, at what rate is oxygen consumed? At what rate is NO produced? At what rate is water vapor produced? 11.18 The following data were obtained in the decomposition of H2O2(aq) to O2(g) and H2O(,). The rate at which oxygen gas was produced was measured. (No oxygen was present initially.)

Time (min)

Volume of O2 (mL)

3.3

10.4

6.9

17.3

30

48

75

30.0

24.9

19.3

13.1

Plot the data and determine the instantaneous rate of decomposition of azomethane at t 5 20 min. Rate Laws and the Concentration Dependence on Rates 11.22 Which of the following would be likely to have a very large rate constant? Which would have a very small rate constant? (a) Dynamite exploding (b) Iron rusting (c) Paper burning 11.23

■ A reaction has the experimental rate equation Rate 5 k[A]2. How will the rate change if the concentration of A is tripled? If the concentration of A is halved?

11.24

■ Second-order rate constants used in modeling atmospheric chemistry are commonly reported in units of cm3 molecule21 s21. Convert the following rate constants to L mol21 s21:

(a) 3.5 3 10214 cm3 molecule21 s21

(b) Calculate the average rate in mL/min for the first 6.9 minutes.

(c) 6.1 3 10230 cm3 molecule21 s21

Time (min)

0

8.3

15.4

19.0

Volume (L)

0

4.2

8.6

11.5

(b) 7.1 3 10218 cm3 molecule21 s21 11.25

11.26

(b) How will the reaction rate change if the NO2 concentration is halved?

Experimental data are listed here for the reaction A : 2 B.

[B] (mol/L)

0.00

0.000

10.0

0.326

20.0

0.572

30.0

0.750

40.0

0.890

(a) Prepare a graph from these data, connect the points with a smooth line, and calculate the rate of change of


The reaction of CO(g) 1 NO2(g) is second-order in NO2 and zero-order in CO at temperatures less than 500 K.

■

(a) Write the rate law for the reaction.

■

Time (s)

■ For each of the rate laws below, what is the order of the reaction with respect to the hypothetical substances X, Y, and Z? What is the overall order?

(a) Rate 5 k[X][Y][Z], (b) Rate 5 k[X]2[Y]1/2[Z], (c) Rate 5 k[X]1.5[Y]21, (d) Rate 5 k[X]/[Y]2

What is the average rate of production of B2 for the first 8.3 min? For the first 19 min?

■

36.2

15

(a) Calculate the average rate in mL/min for the first 3.3 minutes.

11.19 A gas, AB, decomposes and the volume of B2 produced is measured as a function of time. The data obtained are as follows:

11.20

0

(c) How will the reaction rate change if the concentration of CO is doubled? 11.27 Show that if the units of rate are mol L21 s21, then the units of the rate constant in the following second-order reaction are L mol21 s21: H2(g) 1 Br2(g) : 2 HBr 11.28

rate 5 k [H2][Br2]

One reaction that destroys O3 molecules in the stratosphere is

■

NO 1 O3 : NO2 1 O2 When this reaction was studied in the laboratory, it was found to be first order with respect to both NO and O3,


383

with a rate constant of 1.9 3 104 L mol21 s21. If [NO] 5 1.2 3 1025 mol L21 and [O3] 5 2.0 3 1025, what is the rate of this reaction?

Initial [A] Expt. (mol L21)

11.29 The hypothetical reaction, A 1 B : C, has the rate law: Rate 5 k[A]x[B]y

1

0.10

0.10

3.0 3 1024

When [A] is doubled and [B] is held constant, the rate doubles. But the rate increases fourfold when [B] is doubled and [A] is held constant. What are the values of x and y?

2

0.30

0.30

9.0 3 1024

3

0.10

0.30

3.0 3 1024

4

0.20

0.40

6.0 3 1024

11.30 The rate of the decomposition of hydrogen peroxide, H2O2, depends on the concentration of iodide ion present. The rate of decomposition was measured at constant temperature and pressure for various concentrations of H2O2 and of KI. The data appear below. Determine the order of reaction for each substance, write the rate law, and evaluate the rate constant.

Rate (mL min21) 0.090

[H2O2]

11.35

■

0.15

2 NO(g) 1 2 H2(g) : N2(g) 1 2 H2O(g)

0.033

0.178

0.30

0.033

0.184

0.15

0.066

For the reaction

at 1100°C, the following data have been obtained.

[KI]

11.31 Give the order with respect to each reactant and the overall order for the hypothetical reaction

[NO] (mol L21)

[H2] (mol L21)

Rate 5 2D[N2]/Dt (mol L21 s21)

5.0 3 1023

0.32

0.012

1.0 3 10

22

0.32

0.024

1.0 3 10

22

0.64

0.096

Derive a rate law for the reaction and determine the value of the rate constant. 11.36

■

The reaction NO(g) 1 O2(g) : NO2(g) 1 O(g)

A1B1C:D1E which obeys the rate law: Rate 5 k[A][B]2. 11.32 The following experimental data were obtained for the reaction 2A13B:C12D

[A] (mol L21)

[B] (mol L21)

Rate 5 D[C]/Dt (mol L21 s21)

0.127

0.15

0.033

0.127

0.30

0.132

0.255

0.15

0.066

Determine the reaction order for each reactant and the value of the rate constant. 11.33 The following experimental data were obtained for the reaction of NH41 and NO22 in acidic solution. NH41(aq) 1 NO22(aq) : N2(g) 1 2 H2O(,)

[NH41] (mol L21)

[NO22] (mol L21)

Rate 5 D[N2]/Dt (mol L21 s21)

0.0092

0.098

3.33 3 1027

0.0092

0.049

1.66 3 10

27

0.0488

0.196

3.51 3 1026

0.0249

0.196

1.80 3 1026

Determine the rate law for this reaction and calculate the rate constant. 11.34

■ Rate data were obtained at 25°C for the following reaction. What is the rate law expression for this reaction?

A12B:C12D

384

Initial Rate of Initial [B] Formation of C (mol L21) (M min21)


plays a role in the formation of nitrogen dioxide in automobile engines. Suppose that a series of experiments measured the rate of this reaction at 500 K and produced the following data.

[NO] (mol L21)

[O2] (mol L21)

Rate 5 2D[NO]/Dt (mol L21 s21)

0.002

0.005

8.0 3 10217

0.002

0.010

1.6 3 10216

0.006

0.005

2.4 3 10216

Derive a rate law for the reaction and determine the value of the rate constant. 11.37 In a heterogeneous system such as wood burning in oxygen, the surface area of the solid can be a factor in the rate of the reaction. Increased surface area of the wood means increased collisions with oxygen molecules. To understand how surface increases when dividing a solid, do the following. (a) Draw a cube that is 2 cm on a side. What is its total surface area? (b) Divide each face of the cube into four equal sections. If you were able to take the cube apart on the lines you have just drawn, how many cubes would you have? (c) What are the dimensions of the new cubes? (d) What is the total surface area of all the new cubes? What will happen if you divide the cubes again? (e) Now consider burning a log or burning all the toothpicks that could be made from the log. What effect would surface area have on the rate that each burns? Would you rather try to start a campfire with a log or with toothpicks?

■


11.38 In wheat-growing areas, such as the plains of the central United States and Canada, harvested wheat is stored in tall grain elevators that are visible for miles in the flat prairie. The wheat is dumped from trucks into a lower part of the elevator and then moved up the elevators into storage areas. The combustible dust produced as the wheat grains rub together has caused grain elevator explosions. Use the cube analogy in problem 11.37 above to explain why wheat dust dispersed in the air in a grain elevator can react explosively if sparked, whereas whole wheat grains do not.

(c) How long would it take for the concentration of the reactant to drop to 0.040 mol/L? 11.46 A substance undergoes first-order decomposition. After 40.0 min at 500°C, only 12.5% of the original sample remains. What is the half-life of the decomposition? If the original sample weighed 243 g, how much would remain after 2.00 h? 11.47 Show that the half-life of a second-order reaction is given by 1 t1/2 5 — k [A]0

Integrated Rate Laws

In what fundamental way does the half-life of a secondorder reaction differ from that of a first-order reaction?

11.39 The decomposition of N2O5 in solution in carbon tetrachloride is a first-order reaction. 2 N2O5 : 4 NO2 1 O2 The rate constant at a given temperature is found to be 5.25 3 1024 s21. If the initial concentration of N2O5 is 0.200 M, what is its concentration after exactly 10 minutes have passed? 11.40 In Exercise 11.39 above, if the initial concentration of N2O5 is 0.100 M, how long will it take for the concentration to drop to 0.0100 times its original value? 11.41 For a drug to be effective in treating an illness, its levels in the bloodstream must be maintained for a period of time. One way to measure the level of a drug in the body is to measure its rate of appearance in the urine. The rate of excretion of penicillin is first order, with a half-life of about 30 min. If a person receives an injection of 25 mg of penicillin at t 5 0, how much penicillin remains in the body after 3 hours? 11.42 Amoxicillin is an antibiotic packaged as a powder. When it is used to treat babies and small animals, the pharmacist or veterinarian must suspend it in water, so that it can be administered orally with a medicine dropper. The label says to dispose of unused suspension after 14 days. It also points out that refrigeration is required. In the context of this chapter, what is implied in the latter two statements? 11.43 As with any drug, aspirin (acetylsalicylic acid) must remain in the bloodstream long enough to be effective. Assume that the removal of aspirin from the bloodstream into the urine is a first-order reaction, with a half-life of about 3 hours. The instructions on an aspirin bottle say to take 1 or 2 tablets every 4 hours. If a person takes 2 aspirin tablets, how much aspirin remains in the bloodstream when it is time for the second dose? (A standard tablet contains 325 mg of aspirin.) 11.44 A possible reaction for the degradation of the pesticide DDT to a less harmful compound was simulated in the laboratory. The reaction was found to be first order, with k 5 4.0 3 1028 s21 at 25°C. What is the half-life for the degradation of DDT in this experiment, in years? 11.45

If the initial concentration of the reactant in a firstorder reaction A : products is 0.64 mol/L and the halflife is 30.0 s,

■

(a) Calculate the concentration of the reactant exactly 60 s after initiation of the reaction. (b) How long would it take for the concentration of the reactant to drop to one-eighth its initial value?

■


11.48

The following data were collected for the decomposition of N2O5.

■

Time, t (min)

[N2O5] (mol L21)

0

0.200

5

0.171

10

0.146

15

0.125

20

0.106

25

0.0909

30

0.0777

35

0.0664

40

0.0570

(a) Plot ln[N2O5] vs. t and determine the rate constant for this first-order reaction. (b) Find the half-life of the reaction. 11.49 The rate of photodecomposition of the herbicide picloram in aqueous systems was determined by exposure to sunlight for a number of days. One such experiment produced the following data (Hedlund, R.T., Youngson, CR. “The Rates of Photodecomposition of Picloram in Aqueous Systems,” Fate of Organic Pesticides in the Aquatic Environment, Advances in Chemistry Series, #111, American Chemical Society (1972), 1592172):

Exposure Time, t (days)

[Picloram] (mol L21)

0

4.14 3 1026

7

3.70 3 1026

14

3.31 3 1026

21

2.94 3 1026

28

2.61 3 1026

35

2.30 3 1026

42

2.05 3 1026

49

1.82 3 1026

56

1.65 3 1026

Determine the order of reaction, the rate constant, and the half-life for the photodecomposition of picloram.


385

■ The rate of decomposition of SO2Cl2 according to the reaction

SO2Cl2(g) : SO2(g) 1 Cl2(g) can be followed by monitoring the total pressure in the reaction vessel. Consider the following data:

Total Pressure (torr)

0.0

491.7

185.3

549.6

242.8

566.6

304.5

584.1

362.7

599.9

429.5

617.2

509.7

637.0

606.3

659.5

1.0 0.8

What must you do to convert these total pressures into changes in the pressure of the SO2Cl2? Manipulate the data as needed and then use a graphing calculator or spreadsheet to plot the data and determine the order of reaction and the rate constant.

0.6 0.4 0.2 0

Peroxyacetyl nitrate (PAN) has the chemical formula C2H3NO4 and is an important lung irritant in photochemical smog. An experiment to determine the decomposition kinetics of PAN gave the data below. Determine the order of reaction and calculate the rate constant for the decomposition of PAN.

0

100

200

300 Time (s)

400

0

100

200

300 Time (s)

400

300 Time (s)

400

■

Time, t (min)

Partial Pressure of PAN (torr)

0.0

2.00 3 1023

10.0

1.61 3 1023

20.0

1.30 3 10

23

30.0

1.04 3 1023

40.0

8.41 3 10

24

50.0

6.77 3 1024

60.0

5.45 3 1024

11.52 The following series of pictures represents the progress of a reaction in which A2 molecules dissociate into atoms: A2 : 2 A. Each picture represents a “snapshot” of the reaction mixture at the indicated time. t = 0 min

t = 1 min

t = 3 min

500

600

0.0 –0.5 ln[H2O2]

11.51

[H2O2] (mol/L)

Time, t (s)

11.53 Hydrogen peroxide (H2O2) decomposes into water and oxygen. 1 H2O2(aq) : H2O(,) 1 — O2(g) 2 Ordinarily this reaction proceeds rather slowly, but in the presence of some iodide ions (I2), the decomposition is much faster. The decomposition in the presence of iodide was studied at 20°C, and the data were plotted in various ways. Use the graphs below, where concentrations are in moles/liter and time is in seconds, to answer the questions that follow.

–1.0 –1.5 –2.0 –2.5

500

600

12 10 1/ [H2O2] (L/mol)

11.50

8 6 4 2 0

0

100

200

500

600

(a) What is the order of reaction for the decomposition of hydrogen peroxide? Find the rate law for this reaction, including the numerical value of the rate constant.

386


(b) Find the numerical value of the rate constant at 20°C, including the correct units.

■


(c) Obtain an estimate of the initial rate of reaction in the experiment that produced the graphs (i.e., the rate at t 5 0 in the graphs). Temperature and Kinetics 11.54 Use the kinetic-molecular theory to explain why an increase in temperature increases reaction rate. 11.55 The activation energy for the reaction in which CO2 decomposes to CO and a free radical oxygen atom, O, has an activation energy of 460 kJ mol21. The frequency factor is 2 3 1011 s21. What is the rate constant for this reaction at 298 K? 11.56 The labels on most pharmaceuticals state that the medicine should be stored in a cool, dark place. In the context of this chapter, explain why this is sound advice. 11.57 The following rate constants were obtained in an experiment in which the decomposition of gaseous N2O5 was studied as a function of temperature. The products were NO2 and NO3.

k (s21)

Temperature (K)

Allene

9.6 3 105 exp (22750/T )

Propane

6.6 3 106 exp (21970/T )

1-Butene

1.7 3 106 exp (21696/T )

cis-2-Butene

1.9 3 106 exp (2956/T )

(a) Which reaction is likely to be the fastest at 310 K? (b) Which reaction has the lowest activation energy? (c) Which has the highest activation energy? 11.61 Bacteria cause milk to go sour by generating lactic acid. Devise an experiment that could measure the activation energy for the production of lactic acid by bacteria in milk. Describe how your experiment will provide the information you need to determine this value. What assumptions must be made about this reaction? Reaction Mechanisms 11.62 Define the term intermediate in the context of reaction mechanisms.

3.5 3 10

298

2.2 3 10

24

308

6.8 3 1024

11.63 Can a reaction mechanism ever be proven correct? Can it be proven incorrect?

318

11.64

23

328

Determine Ea for this reaction in kJ/mol. 11.58 The table below presents measured rate constants for the reaction of NO with ozone at three temperatures. From these data, determine the activation energy of the reaction in kJ/mol.

k (L mol21 s21)

Temperature (K)

1.3 3 106

200

4.4 3 10

6

250

9.9 3 106

300

11.59 The table below presents rate constants measured at three temperatures for the following reaction, which is involved in the production of nitrogen oxides in internal combustion engines. O(g) 1 N2(g) : NO(g) 1 N(g)

■

Arrhenius Expression

25

3.1 3 10

11.60

Volatile Organic Chemical

k (L mol21 s21)

Temperature (K)

4.4 3 102

2000

2.5 3 105

3000

5.9 3 106

4000

What is the rate law for each of the following elementary reactions?

■

(a) Cl(g) 1 ICl(g) : I(g) 1 Cl2(g) (b) O(g) 1 O3(g) : 2 O2(g) (c) 2 NO2(g) : N2O4(g) 11.65 Describe how the Chapman cycle is a reaction mechanism. What is the molecularity of each reaction in the Chapman cycle? 11.66 What are the reaction intermediates in the Chapman cycle? 11.67 Propane is used as a household fuel in areas where natural gas is not available. The combustion of propane is given by the equation C3H8 1 5 O2 : 3 CO2 1 4 H2O Does this equation represent a probable mechanism for the reaction? Defend your answer. 11.68 Five people are responsible for cleaning up after a very large banquet. Two are clearing the tables, while one washes the dishes and two dry. Which step is likely to be rate-determining? 11.69 The following mechanism is proposed for a reaction: NO 1 Br2 : NOBr2

(slow)

NOBr2 1 NO : 2 NOBr

(fast)

(a) Write the overall equation for the reaction.

Determine the activation energy of the reaction in kJ/mol.

(b) What is the rate-determining step?

■ Many reactions occur in the formation of photochemical smog, including the reaction of ozone with various volatile organic chemicals, or VOCs. The table in the next column shows the Arrhenius expression for four such reactions.

(d) What is the molecularity of each step of the reaction?


(c) What is the intermediate in this reaction? (e) Write the rate expression for each step. 11.70 HBr is oxidized in the following reaction: 4 HBr(g) 1 O2(g) : 2 H2O(g) 1 2 Br2(g)


387

A proposed mechanism is HBr 1 O2 : HOOBr

(slow)

HOOBr 1 HBr : 2 HOBr

(fast)

HOBr 1 HBr : H2O 1 Br2

(fast)

(a) Show that this mechanism can account for the correct stoichiometry.

(a) Write the equation for the decomposition of H2O2.

(b) Identify all intermediates in this mechanism.

(b) What role does MnO2 play?

(c) What is the molecularity of each elementary step?

(c) In the chemistry lab, a student substituted a chunk of MnO2 for the powdered compound. The reaction rate was not appreciably increased. What is one possible explanation for this observation?

(d) Write the rate expression for each elementary step. (e) Identify the rate-determining step. 11.71

decomposes slowly over time, and the rate of decomposition increases with an increase in temperature and in the presence of light. However, the rate of decomposition increases dramatically if a small amount of powdered MnO2 is added to the solution. The decomposition products are H2O and O2. MnO2 is not consumed in the reaction.

■ The reaction of NO2(g) and CO(g) is thought to occur in two steps:

Step 1: NO2(g) 1 NO2(g) : NO(g) 1 NO3(g)

(slow)

Step 2: NO3(g) 1 CO(g) : NO2(g) 1 CO2(g)

(fast)

(a) Show that the elementary steps add up to give the overall stoichiometric equation. (b) What is the molecularity of each step? (c) For this mechanism to be consistent with kinetic data, what must be the experimental rate equation? (d) Identify any intermediates in this reaction.

(d) Is MnO2 part of the stoichiometry of the decomposition of H2O2? INSIGHT INTO Tropospheric Ozone 11.81 Why is smog appropriately labeled photochemical smog? 11.82 When oxygen molecules, O2, react with oxygen atoms to form ozone, a third molecule must be present. (a) What is the molecularity of this elementary reaction? (b) Would you expect this reaction to be fast or slow? Explain your reasoning. (c) Why is the third molecule required?

Catalysis 11.72 If a textbook defined a catalyst as “a substance that increases the rate of a reaction,” would that definition be adequate? 11.73 The word catalyst has been adopted outside of science. In sports, a reserve who enters the game might be described as the catalyst for a comeback victory. Discuss how this use of the word is similar to the scientific use of the word. How is it different from the scientific use of the word? 11.74 In the process by which CFCs contribute to the depletion of stratospheric ozone, is light a catalyst? Why or why not? 11.75 What distinguishes homogeneous and heterogeneous catalysis? 11.76 Define the terms adsorbed and desorbed. Do the processes to which these terms refer occur only in heterogeneous catalysis? 11.77 Based on the kinetic theory of matter, what would the action of a catalyst do to a reaction that is the reverse of some reaction that we say is catalyzed? 11.78 Draw a hypothetical activation energy diagram for an uncatalyzed reaction. On the same set of coordinates, draw a similar graph for the catalyzed reaction. Why does reaction rate increase with the use of the catalyst?

11.83 What is a VOC? What role do VOCs play in photochemical smog? 11.84 Why does the formation of ozone in smog depend on having less NO(g) present? 11.85 Many urban centers in North America do not meet air quality standards for ozone levels. What factors associated with living in a city contribute most strongly to the formation of potentially dangerous levels of ozone? 11.86 Ozone is not stable at pressures typically found in the troposphere. How might this contribute to the variation of ozone levels over the course of each day? 11.87 Why do ozone concentrations lag in time relative to other pollutants in photochemical smog? 11.88 Reducing ozone formation in photochemical smog would likely require action at both the chemical and societal levels. (a) Propose strategies at the chemical level—i.e., which chemicals would need to be accentuated or eliminated? (b) Propose strategies at the societal level—i.e., what changes would the average person need to make to accomplish the changes you suggested at the chemical level? 11.89 Find a website that tells you whether you live in an area that has unhealthy levels of ozone. (a) What is the URL of the site? Does this site strike you as a reliable source of information?

11.79 The catalytic converter of your automobile is likely to contain two different catalysts. Based on the fact that the role of this device is to convert NOx to N2 and O2 as well as to convert CO to CO2, discuss why it makes sense that there are two catalysts. (HINT: Consider the oxidationreduction nature of the reactions involved.)

11.90 Use the web to find out what it means if the U.S. Environmental Protection Agency lists a city as a nonattainment area.

11.80 The label on a bottle of 3% (by volume) hydrogen peroxide, H 2O 2, purchased at a grocery store, states that the solution should be stored in a cool, dark place. H2O2

11.91 Western Michigan can experience higher than normal ozone levels in the summer because prevailing southwest winds carry industrial pollution in the form of VOCs from

388


(b) Does your area have unhealthy air?

■


Milwaukee and Chicago. (For maps, visit the EPA Region 5 website at http://www.epa.gov/region5) (a) Rewrite equations 11.10.d2f, using hexane, C6H14, as the VOC pollutant. (b) Explain how the NO2 produced in the reactions you wrote in (a) can result in tropospheric ozone production in Western Michigan. 11.92 Along the Front Range of Colorado, the greatest contributor to Denver’s “brown cloud” is vehicle traffic. Motor vehicles emit VOCs while they stir up the dust on the highways, increasing particulates. (Visit the Regional Air Quality Control’s website at http://www.raqc.org/ for more information.) (a) In 2007, Denver exceeded the EPA’s 8-hour ozone standard, which at that time was 0.08 ppm. Imagine that you are responsible for helping Denver (or your community) decrease ozone concentrations. Using the information in Table 11.3, list four things you would tell Denver residents to do around the house that would decrease emissions of VOCs and thus decrease ozone formation.

pressures. Discuss the differences in reaction conditions for these three reactions in terms of activation energy and collision theory. 11.95

On a particular day, the ozone level in Milwaukee exceeded the EPA’s 1-hour standard of 0.12 ppm by 10 ppb. How many ozone molecules would be present in 1 liter of air at the detection site?

11.96

The EPA’s 8-hour ozone standard for communities is 0.075 ppm. On a certain day, the concentration of ozone in the atmosphere in Denver was 1.5 3 10 23 mol/L. Was Denver in compliance with EPA regulations that day?

11.97

The following is a thought experiment. Imagine that you put a little water in a test tube and add some NaF crystals. Immediately after you add NaF, you observe that the crystals begin dissolving. The quantity of solid NaF decreases, but before long, it appears that no more NaF is dissolving. The solution is saturated. (a) The equation for the dissolution of NaF in water is NaF(s) : Na1(aq) 1 F2(aq). As NaF dissolves, what do you think happens to the rate of dissolution? Describe what is occurring on the molecular level.

(b) One of the suggestions given to residents of the Denver area to help decrease ozone formation is to fill their car’s gas tank after dark. Explain why this could be effective.

(b) Assume that the reverse reaction, Na1(aq) 1 F2(aq) : NaF(s), also occurs as the crystal dissolves. In other words, both dissolution and precipitation are taking place. When it appears that there is no more change in the quantity of NaF dissolving (the solution is saturated), what has happened to the rates of the forward and reverse reactions? Explain your answer.

Additional Problems 11.93 According to an article published in 1966 in the Journal of Chemical Education, xenon reacts photochemically with fluorine at room temperature (298 K) to produce XeF2. A portion of a table showing ratios of gases in the reaction vessel for four experiments is reproduced below. (Assume that atmospheric pressure is exactly 1 atm.) The reaction is as shown.

11.98

H2(g) 1 I2(g) : 2 HI(g)

Xe(g) 1 F2(g) : XeF2(s)

Experiment Volume of Xe, in mL at atmospheric pressure Volume of F2, in mL at atmospheric pressure

1

2

3

Rate 5 k[H2][I2]

Determine which of the following statements are true. If a statement is false, indicate why it is incorrect.

4

(a) The reaction must occur in a single step. 300. 271

350. 260

400. 262

450.

(b) This is a second-order reaction overall.

183

(c) Raising the temperature will cause the value of k to decrease. (d) Raising the temperature lowers the activation energy for this reaction.

(Data from John H. Holloway, J. Chem. Ed., 43, 203 (1966).)

(a) Draw the Lewis structure of XeF 2 and predict the molecule’s shape.

(e) If the concentrations of both reactants are doubled, the rate will double.

(b) Calculate the partial pressure of each gas in Experiment 3 if the amounts shown are combined such that the total pressure remains 1 atm.

(f) Adding a catalyst in the reaction will cause the initial rate to increase.

(c) Calculate the theoretical yield of XeF2 in Experiment 3. (d) After the reactants were introduced, the bulb was placed in sunlight. XeF2 began crystallizing on the inner surface of the bulb about 24 hours later. Compare the rate of this reaction to the rate of precipitation of CaF2 from solutions of Ca21(aq) and F2(aq) in the laboratory. 11.94 In the previous problem, the reaction to produce XeF2 occurs photochemically at room temperature. However, to produce XeF 4(s), higher temperature and pressure are required. Production of XeF6 requires even higher

■

The following statements relate to the reaction for the formation of HI:

■


11.99

The Haber process, which produces ammonia from nitrogen and hydrogen, is one of the world’s most important industrial chemical reactions. N2(g) 1 3 H2(g) : 2 NH3(g) For this reaction, both DH ° and DS° are negative, meaning that product formation is thermodynamically favored by low temperatures. Yet the industrial reaction is commonly carried out at temperatures between 4002600°C. These high temperatures require a significant input of energy; some estimates indicate that the industrial production of ammonia is responsible for 1% of the world’s


389

11.100

total energy consumption. Suggest a reason why these high temperatures are employed.

tablish the age of the artifact? What would you have to know or look up?

Experiments show that the reaction of nitrogen dioxide with fluorine

11.106 You are told that an artifact is 3200 years old. How would you use this information to predict the number of radiation counts from carbon-14 that you would measure?

■

2 NO2(g)1F2(g) : 2 FNO2(g) has the rate law Initial reaction rate 5 k[NO2][F2] The reaction is thought to occur in two steps. Step 1: NO2(g) 1 F2(g) : FNO2(g) 1 F(g) Step 2: NO2(g) 1 F(g) : FNO2(g) (a) Show that the sum of this sequence of reactions gives the balanced equation for the overall reaction. (b) Which step is rate determining? 11.101

Substances that poison a catalyst pose a major concern for many engineering designs, including catalytic converters. One design concept is to add materials that react with potential poisons before they reach the catalyst. Among the commonly encountered catalyst poisons are silicon and phosphorus, which typically form phosphate or silicate ions in the oxidizing environment of an engine. Group 2 elements are added to the catalyst to react with these contaminants before they reach the working portion of the catalytic converter. If estimates show that a catalytic converter will be exposed to 625 g of silicon during its lifetime, what mass of beryllium would need to be included in the design?

■

11.102 Chemical engineering design often uses mol% as a measure of concentration. A combustion reactor has an inlet flow of 5.50 mol% propane in air. Analysis of the product stream from the reactor finds 6.25 mol% CO2. Does this measurement indicate complete combustion? What other species might you suggest monitoring in order to determine if the reaction kinetics are working in a predictable way. 11.103

■ The flashing of fireflies is the result of a chemical reaction, and the rate of flashing can be described by the Arrhenius equation. A certain batch of fireflies were observed to flash at a rate of 16.0 times per minute at 25°C and at a rate of 5.5 times per minute at 15°C. Use these data to find the apparent activation energy for the reaction that causes the flies to flash.

FOCUS ON PROBLEM SOLVING EXERCISES 11.104 If you have a reaction for which you do not know the order, explain how you can obtain the information you need with only a single experiment. 11.105 You measure an archeological specimen and find that the count rate from carbon-14 is x% of that seen in living organisms. How does this information allow you to es-

390


11.107 You are designing a self-heating food package. The chemicals you are using seem to be used up too quickly. What experiments could be run to devise conditions that would allow the heating pack to last longer? Cumulative Problems 11.108 Use the web to look up the rate of reaction for alkali metals with water. How can you explain the trend in these rates in terms of the concepts presented in this chapter? 11.109 Fluorine often reacts explosively. What does this fact suggest about fluorine reactions at the molecular perspective? 11.110 In Chapter 4, we noted that compressibility is an important factor for fuels, and that some hydrocarbons are more compressible than others. If a fuel is not very compressible, what can we infer about the activation energy for its combustion reaction? 11.111 A free radical is often used as a catalyst for polymerization reactions. Such reactions normally need to be carried out at higher temperatures. Explain this fact in terms of the molecular perspective and the activation energy of the reactants. 11.112

When formic acid is heated, it decomposes to hydrogen and carbon dioxide in a first-order decay. ■

HCOOH(g) : CO2(g) 1 H2(g) The rate of reaction is monitored by measuring the total pressure in the reaction container.

Time (s)

P (torr)

0

220

50

324

100

379

150

408

200

423

250

431

300

435

Calculate the rate constant and half-life in seconds for the reaction. At the start of the reaction (time 5 0), only formic acid is present. (HINT: Find the partial pressure of formic acid using Dalton’s law of partial pressure and the reaction stoichiometry to find PHCOOH at each time.)

■


Chemical Equilibrium

12 OUTLINE 12.1 INSIGHT INTO Concrete Production and Weathering 12.2 Chemical Equilibrium 12.3 Equilibrium Constants 12.4 Equilibrium Concentrations 12.5 LeChatelier’s Principle 12.6 Solubility Equilibria 12.7 Acids and Bases 12.8 Free Energy and Chemical Equilibrium 12.9 INSIGHT INTO Borates and Boric Acid

The Natchez Trace Parkway Bridge, in Williamson County, Tennessee, is a concrete double arch bridge. The unique open design of the bridge places strict strength requirements on the concrete used. The weight of the bridge is concentrated at the crown of each arch. National Park Service

I

n our study of kinetics, we saw that the rate of a reaction changes over time, invariably slowing down as the reaction progresses. If we were to observe many reactions, this behavior would be evident each time. What does this imply? Eventually, the net rate of any chemical reaction approaches zero. When the rate goes to zero, the reaction has reached a state of equilibrium. The concept of chemical equilibrium is critically important, and we will briefly look at several aspects of equilibrium in this chapter. Let’s begin by considering a very familiar building material—concrete.

Chapter Objectives After mastering this chapter, you should be able to ❚ list chemical reactions important in the production and weathering of concrete. ❚ explain that equilibrium is dynamic and that at equilibrium, the forward and backward reaction rates are equal. State these ideas in your own words. ❚ write the equilibrium constant expression for any reversible reaction. ❚ calculate equilibrium constants from experimental data. ❚ calculate equilibrium composition from initial data and the numerical value of the equilibrium constant. ❚ calculate molar solubility from Ksp or vice versa.


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❚ write equilibrium constants for the dissociation of weak acids and weak bases and use them to calculate pH or the degree of ionization. ❚ use LeChatelier’s principle to explain the response of an equilibrium system to applied stresses. ❚ calculate the new equilibrium composition of a system after an applied stress. ❚ explain the importance of both kinetic and equilibrium considerations in the design of industrial chemical processes.

INSIGHT INTO

12.1

The name Portland cement was first used in an 1824 patent in England by Joseph Aspdin who thought the material looked like a type of limestone quarried near Portland, England.

Concrete Production and Weathering

There are few building materials that are more ubiquitous than concrete. When people think about concrete, usually they associate it with traits like durability and strength. You may never have associated concrete with chemistry, but there are actually a number of complex chemical systems that must be considered in many engineering designs that use concrete. Some of the most interesting chemistry occurs in the weathering of concrete due to prolonged exposure to the environment. Before we can consider weathering of concrete, we will need to look at its production. Traditionally concrete has been composed of cement, water, and aggregate. Modern concrete preparation includes the use of additional components called admixtures to help manipulate the concrete into having desired properties, either as it is being poured or in long-term usage. From a chemical perspective, the most interesting reactions arise in the preparation of the cement and the admixtures. Most concrete uses Portland cement. Portland cement starts with the production of CaO from limestone, which is mostly calcium carbonate, CaCO3. CaCO3 : CaO 1 CO2 This step in the production of cement for concrete accounts for an estimated 5% of CO2 released into the atmosphere annually. In addition to CaO, cement includes oxides of silicon and aluminum. These combined materials are then hydrated (have water added) when the concrete is mixed. The resulting mixture may have a fairly wide range of composition, so the equations that represent the chemistry are often written to include a variable, x. This variable may assume different values depending on the composition of a particular concrete but doesn’t fundamentally alter the nature of the chemical reaction. Three representative reactions for the hydration process are: 3 CaO?Al2O3 1 6 H2O : Ca3Al2(OH)12 2 CaO?SiO2 1 x H2O : Ca2SiO4?xH2O 3 CaO 1 SiO2 1 (x 1 1)H2O : Ca2SiO4?xH2O 1 Ca(OH)2 If you count the chemical bonds in the reactants and products of these reactions, it should not be surprising that they release energy as heat, because the net result of the process is the formation of additional chemical bonds. Figure 12.1 shows a graph of the energy liberated by the hydration of concrete as a function of time. Note that the reaction is fairly slow: the time scale in the graph is more than a month. This helps explain why engineers who use concrete need to specify curing times to achieve the desired characteristics of the concrete for their designs. Recently, the use of fly ash to partially replace Portland cement has become increasingly common. Fly ash is generated when coal is burned in power plants. Minerals that are present in the coal react with oxygen at high temperatures and produce a residue known as fly ash. The average composition of fly ash is similar to that of Portland cement, with the four main components being SiO2, Al2O3, Fe2O3, and CaO.

392

Chapter 12 Chemical Equilibrium

Figure 12.1 ❚ The hydration

Heat of hydration (kJ/mol )

50

reaction for concrete formation is exothermic and takes place over relatively long times. After a rapid temperature rise, called the preinduction period, the main energy release occurs about a week after the original mixing of the concrete.

40

30

20

10

0

6

12

18 24 30 Hydration time (days)

36

42

48

Fly ash typically consists of small spherical particles, and the incorporation of these particles can improve the strength of concrete. Inclusion of fly ash in concrete also provides a use for a material that might otherwise be considered waste. The use of fly ash is increasingly suggested as a more environmentally friendly way to manufacture concrete cement. From the perspective of design considerations, modern engineers have notably more options in the specifications of concrete because of the use of admixtures. Table 12.1 indicates uses for admixtures and common chemicals that provide the desired characteristics. Water reducers are used to lower the amount of water in the concrete while not reducing the ability to work with the material. Air entraining admixtures improve durability of cement by stabilizing small air bubbles within the cement portion of the concrete, particularly when concrete is exposed to water during freeze-thaw cycles. Other forms of admixtures combat the effects of moisture via waterproofing. Finally, some admixtures are useful as either accelerators of the hardening process Table

❚

12.1

Chemical composition of selected admixtures

Function

Compound

Origin

Water reduction

Lignosulfonate

Wood/pulp byproduct

Water reduction

Hydroxycarboxylic acids

Chemical production

Air entrainment

Abietic and pimeric acids

Wood resins

Air entrainment

Alkyl-aryl sulphonates

Industrial detergents

Waterproofing

Fatty acids

Vegetable and animal fats

Acceleration

Calcium chloride

Chemical production

Acceleration

Calcium formate

Chemical production byproduct

Acceleration

Triethanolamine

Chemical production

Retardation

Borates

Borax mining

Retardation

Magnesium salts

Chemical production

12.1 Concrete Production and Weathering

393

Dr. Peter C. Taylor

Figure 12.2 ❚ Treatment of concrete with phenolphthalein reveals the chemical effects of weathering. The pink color shows that the bulk of this sample remains basic, even after extended exposure to the elements. The colorless area at the top, which is the surface that was actually exposed, is no longer basic. Reaction with atmospheric CO2 has neutralized the hydroxide in this surface region. Most retarders also serve as waterreducing admixtures.

or retarders of that process. Once in the concrete, any of these admixtures can play a role in how it functions in the environment. Weathering of concrete has several components, including the freeze-thaw cycle just noted as being partially mitigated by air entraining admixtures. We will focus on those components whose effects are primarily chemical from the start. A key chemical reaction related to the aging of concrete is carbonation. In carbonation, CO2 from the air diffuses into the concrete. Once there, it can react with calcium hydroxide present in essentially a two-step process represented by the following reactions: Ca(OH)2(s) : Ca21(aq) 1 2 OH2(aq) Ca21(aq) 1 2 OH2(aq) 1 CO2(g) : CaCO3(s) 1 H2O(,)

We introduced the concepts of acids and bases in Section 3.3 and will explore them further in this chapter.

As we have noted earlier, the presence of hydroxide ions defines a base. So it is possible to observe the carbonation of concrete using an acid-base indicator. Figure 12.2 shows a cross section of a piece of concrete that has been treated with an indicator called phenolphthalein, which turns pinks in the presence of a base. The bulk of this sample appears pink, confirming that the interior of the concrete remains basic. But the top 1/8 inch, which is the surface that was actually exposed, is not colored at all. This indicates that CO2 from the environment has reacted with and neutralized the hydroxide originally present in the concrete. As is true in the environmental exposure of most engineering designs, the reactions that take place in concrete have long times in which to occur, and as a result, they are capable of reaching chemical equilibrium.

12.2

Chemical Equilibrium

Equilibria in applied chemical systems such as concrete are unavoidably complex. But we can still understand much of the chemistry of these complex equilibria in terms of a few underlying principles. So as we begin our exploration of equilibria, we will use simpler systems to allow us to focus on those principles.

Forward and Reverse Reactions If you place a glass of water on a table and leave it overnight, in the morning you may find an empty glass, or at least find that the water level has decreased (Figure 12.3). This won’t surprise you because you know that the missing water has evaporated.

394



Figure 12.3 ❚ This photo sequence shows the water level in two glasses over the course of time. The glass on the left is covered, whereas that on the right is open to the air. In the left-hand photo, the glasses are filled to the same level. In the second photo, taken 10 days later, the water level in the right-hand glass is visibly lower, whereas that in the left-hand glass remains unchanged. The trend continues in the final photo.

Try the experiment again, this time covering the glass, and you’ll find that all of the water stays in the glass. In a closed system—like the covered glass—a dynamic equilibrium is established between liquid water and water vapor (Figure 12.4). At equilibrium, the rate at which water molecules leave the surface of the liquid equals the rate at which they return to the surface. The rate of evaporation will be determined by the distribution of molecular energies in the liquid. Only those molecules with sufficient kinetic energy can overcome the intermolecular forces in the liquid and escape into the vapor phase. Assuming that the water is at a constant temperature, the rate of evaporation will be constant. What about the rate of condensation? Each time a gas phase molecule strikes the surface of the liquid, it has some probability of sticking to

Rate of evaporation

Rate of condensation

(a)

(b)

(c)

Figure 12.4 ❚ The equilibrium between liquid and vapor in a closed container is governed by the kinetics of evaporation and condensation. Here a liquid is placed into a closed container with no vapor phase present. Energetic molecules at the surface of the liquid can move into the gas phase, as shown in (a). Once some vapor is present, gas molecules will occasionally strike and stick to the liquid surface, so condensation begins to compete with evaporation, as seen in (b). The rate of condensation rises over time, whereas the rate of evaporation stays constant. Eventually the two rates become equal. In this equilibrium state, both evaporation and condensation continue, as shown in (c).

12.2 Chemical Equilibrium

395

We introduced the idea of dynamic equilibrium when discussing vapor pressure in Section 8.5.

that surface and returning to the liquid phase. The rate at which gas molecules strike the surface of the liquid is proportional to the number of gas molecules present, or to the gas pressure. If only liquid is present initially, the rate of condensation will be zero. But because evaporation is occurring, after a short time there will be some vapor in the container. That means condensation will begin. As long as the (constant) rate of evaporation is greater than the rate of condensation, the net effect will be to increase the pressure of the vapor phase. That, in turn, will increase the rate of condensation. Eventually the two rates will become equal, and after that point, no net change in the amounts of liquid or vapor will be seen. We call this a dynamic equilibrium because the evaporation and condensation do not stop at equilibrium. The rates of these competing processes become equal to one another, but they do not go to zero. So on the microscopic level, individual molecules continue to move from the liquid to the vapor phase and back again. But from a macroscopic perspective, the observable amounts of liquid and vapor no longer change. Turning to chemical reactions, the manner in which we have written chemical equations thus far has been to consider only the forward direction of the reaction. In the reverse reaction, the substances we’ve been calling products react to form reactants. In some cases, the reverse reaction may occur to a large degree and in other cases to an infinitesimal and immeasurable degree, but in principle, it can occur in every chemical reaction. This means that in a closed system, any chemical reaction will eventually reach a dynamic equilibrium analogous to that we just described for evaporation. At the instant we initiate a chemical reaction between two reactants, there are no products present (Figure 12.5). We have already seen that the rate of a chemical reaction depends on the concentrations of reactants, so initially the rate of the reverse reaction must be zero because the concentrations of its reactants—which are the products of the forward reaction—are zero. Thus when we first mix the reactants, the rate of the forward reaction is greater than that of the reverse reaction. Over time, the concentrations of the reactants decrease, and those of the products increase. These changes in concentration are accompanied by changes in rate: the forward reaction slows down and the reverse reaction speeds up. Ultimately, chemical equilibrium is reached when the two rates become equal, and the observable concentrations of both reactants and products become constant. In any chemical system at equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

Reactants

Products

Concentration

Reactants

Products

This graph implies that the reaction does not proceed very far toward products. We’ll see next how this type of observation can be quantified. Time

Figure 12.5 ❚ The graph depicts the approach of a chemical system to equilibrium. Initially only reactants are present; the forward reaction proceeds and forms products. As the product concentrations increase, the reverse reaction becomes increasingly important. Eventually the rates of the forward and reverse reactions become equal, and the measured concentrations of reactants and products all become constant at equilibrium.

396


Mathematical Relationships The rate laws of the forward and reverse reactions suggest that a mathematical relationship can be written to describe equilibrium. For any reaction involving reactants R and products P, we write the chemical equation at equilibrium with two arrows, one pointing in each direction, to emphasize the dynamic character of the process: RNP For simplicity, we will assume that this reaction proceeds in each direction through a single elementary step. (See Section 11.6. The result of our derivation here does not actually depend on this assumption, but it is easier to understand this way.) We can write the rate law for the forward direction of this reaction as: Ratefor 5 kfor[R] Similarly, for the reverse reaction, we can write: Raterev 5 krev[P] We know that these two rates must be equal at equilibrium, so Ratefor 5 Raterev Therefore, kfor[R]eq 5 krev[P]eq Because the rates are equal only if the system is at equilibrium, we have explicitly designated the concentrations as equilibrium concentrations here by adding a subscript “eq.” We can manipulate this equation to obtain kfor krev

[P]eq

—5 —

[R]eq

Both kfor and krev are rate constants, so as long as temperature is constant, the lefthand side of this equation is also constant. This means that at a given temperature, the ratio [P]eq/[R]eq is a constant. (Remember that rate constants vary with temperature, so this ratio may also change as a function of temperature.) This expression will be at the heart of our mathematical treatment of equilibrium in Section 12.3. To make this idea more concrete, consider what happens when salicylic acid, C6H4(OH)COOH, a water-reducing admixture component, dissolves in water. O

OH OH

salicylic acid

As we will see in greater detail in Section 12.7, an acid dissolves in water to form hydrogen ions and an anion, in this case C6H4(OH)COO2: C6H4(OH)COOH (aq) N H1(aq) 1 C6H4(OH)COO2(aq) If we assume that the reactions in both directions are elementary rate processes, the forward and reverse rate laws can be written: Ratefor 5 kfor[C6H4(OH)COOH] Raterev 5 krev[ H1][C6H4(OH)COO2] Setting these rates equal to one another at equilibrium gives us kfor[C6H4(OH)COOH] 5 krev[ H1][C6H4(OH)COO2]

12.2 Chemical Equilibrium

397

We can rearrange this so that the two rate constants are on one side and the concentrations are on the other: kfor krev

c

H1 d eq cC6H4 (OH)COO2d eq

— 5 —— 5 constant

12.3

c

C6H4(OH)COOHd eq

Equilibrium Constants

We noted earlier that Portland cement includes calcium oxide that is mixed with other materials. The decomposition of limestone, CaCO3, noted in Section 12.1, occurs in a kiln via the reaction: CaCO3(s) : CaO(s) 1 CO2(g) The extent of this reaction and therefore the amount of CaO recovered for use in the cement is an example of an equilibrium process. For this reaction to be used in an industrial process such as the manufacture of cement, it is important to know how much CaO is available, and because of the environmental concerns associated with increased atmospheric CO2 level, the amount of this gas that is released. The need for quantitative answers to questions like this requires a mathematical model for equilibrium. In the previous section, we saw that a relationship exists between concentrations of reactants and products. This relationship, which was first proposed by Cato Maximilian Guldberg and Peter Waage in 1864, is often called the law of mass action.

The Equilibrium (Mass Action) Expression For a general chemical reaction, aA1bBNcC1dD we can define a ratio of concentrations, whether or not at equilibrium: [C]c[D]d Q 5— [A]a[B]b

(12.1)

We refer to this ratio as the reaction quotient, Q. Product concentrations appear in the numerator of the expression, and reactants appear in the denominator. Each concentration is raised to a power corresponding to the stoichiometric coefficient of that species in the balanced chemical equation for the equilibrium. At equilibrium, however, this ratio becomes the equilibrium expression, and the corresponding value of Q is called the equilibrium constant, K. Here we use the subscript “eq” to emphasize that the system is at equilibrium. This subscript is often omitted, and the use of K rather than Q implies equilibrium.

[C]ceq[D]deq K5— [A]aeq[B]beq

(12.2)

To determine the value of K for a particular reaction, we must measure the concentrations of the reactants and products at equilibrium. Note that, unlike rate expressions in chemical kinetics, the expression for the equilibrium constant is always based directly on the reaction stoichiometry. The following Example Problem shows how to write an equilibrium expression.

E XAM PL E PRO BLEM

12 .1

Concrete used in smokestacks has to be designed to withstand sometimes caustic conditions. Sulfur oxides in particular are produced in some cases, and they would

398


establish equilibrium if they did not disperse: 2 SO2(g) 1 O2(g) N 2 SO3(g) Write the equilibrium expression for this reaction.

Strategy We use the definition of the equilibrium expression. Write the products in the numerator and the reactants in the denominator. The exponent for each concentration term is given by the corresponding stoichiometric coefficient from the chemical equation.

Solution c

SO3d 2

K5— 2 cSO d cO d 2 2

Analyze Your Answer The only way we can assess an answer like this is to make sure that we used the definition of the equilibrium constant correctly. We have the product (SO3) in the numerator and the reactants (SO2 and O2) in the denominator, which is correct, and we have exponents to match the stoichiometric coefficients. So our answer is correct.

Check Your Understanding Nitrogen oxides in the atmosphere may acidify rainwater. Write the equilibrium expression for the following reaction in which gaseous nitrous acid forms in the atmosphere: NO(g) 1 NO2(g) 1 H2O(g) N 2 HNO2(g)

Gas Phase Equilibria: Kp vs. Kc Many equilibrium reactions, including that in the previous example, take place in the gas phase. The equilibrium expression we wrote in Example Problem 12.1 includes the molar concentrations of four gases. Although there is nothing wrong with such concentrations, it is often easier to describe gases using partial pressures rather than molar concentrations. Hence an alternative definition of the equilibrium constant in which partial pressures replace molarities is sometimes used for gas phase reactions. The symbol Kp is used to indicate such an equilibrium constant. Let’s return to the general equilibrium we used to develop Equations 12.1 and 12.2 and modify it slightly to show all species as gases: a A(g) 1 b B(g) N c C(g) 1 d D(g) From this we can write an expression for Kp: (PC)ceq (PD)deq Kp 5 — (PA)aeq (PB)beq

(12.3)

Because we now have two alternative definitions for equilibrium constants, it is appropriate to ask whether they will lead to the same values. To answer that question, we need to relate the molar concentration of a gas to its partial pressure. That connection can be made through the ideal gas law and more specifically by Equation 5.5 (page 137), which defines the partial pressure of a gas i: niRT Pi 5 — V The molar concentration of this gas is just (ni/V ). Switching from “i” to “A” to match our generalized reaction, we can write nA PA 5 — RT 5 [A]RT V

12.3


399

Similar expressions can be written for the remaining gases in our reaction (B, C, and D). Inserting all of these into Equation 12.3, ([C]eqRT )c ([D]eqRT )d Kp 5 —— ([A]eqRT )a ([B]eqRT)b [C]ceq[D]deq 5— 3 RT (c 1 d 2 a 2 b) [A]aeq[B]beq 5 Kc 3 RT (c 1 d 2 a 2 b) Here the subscript ‘c’ on the last K is added to emphasize that this is the equilibrium constant written in molar concentrations. It is immediately clear that unless RT (c 1 d 2 a 2 b) is equal to one, Kp and Kc will not be equal. The exponent (c 1 d 2 a 2 b) in that term represents the change in the number of moles of gas from the reactants to the products. Analogously to the way we have written the changes in various thermodynamic quantities, we can write this as Dngas, defined as the number of moles of gas on the product side minus the number of moles of gas on the reactant side. Using this definition, we can write Kp 5 Kc 3 RT (Dngas)

(12.4)

The two equilibrium constants, Kp and Kc, will be equal only if Dngas 5 0. When working with gas phase equilibrium constants, therefore, it is always important to be certain that you know whether the value you look up is Kp or Kc. For simplicity, all equilibrium constants in this text will be based on molar concentrations, not pressures. Accordingly, we will not use the subscript ‘c’ when writing these K values.

Homogeneous and Heterogeneous Equilibria The equilibria we have examined so far are all examples of homogeneous equilibria in which reactants and products are in the same phase, either gaseous or aqueous. If any of the reactants or products is in the solid or liquid phase, though, we will need to adjust our formulation of the equilibrium expression slightly. The concentration of a pure solid or liquid will not change as the reaction proceeds. Indeed, the phrase “concentration of a pure solid or liquid” should strike you as a bit odd; unlike gases or solutions, solids and liquids do not have variable concentrations. For example, suppose that the decomposition of calcium carbonate that occurs in the making of Portland cement were carried out in a closed system (Figure 12.6): CaCO3(s) N CaO(s) 1 CO2(g) Once equilibrium has been established at a particular temperature, the concentration of gaseous CO2 can be expressed in mol/L. If the system is heated, more CaO and CO2 form, and, if the volume of the system is held constant, the concentration of CO2 in mol/L increases as equilibrium is established at a higher temperature. But the concentrations of the solids do not change because, as the numbers of moles of each changes, so do their volumes. The densities of solids change very little with a change in temperature, and the ratio of the number of moles (and mass) of CaCO3, for example, to the volume of CaCO3 remains constant. A solid’s concentration does not change, nor does that of a liquid. This fact has led to a different way to describe the equilibrium expression for heterogeneous equilibrium in a multiphase system. Because their concentrations are constant, reactants and products that are either liquid or solid do not appear in the expression, so the equilibrium expression for the system above includes only the carbon dioxide concentration. K 5 [CO2] In the following example, you can see how this idea can be applied to precipitation reactions.

400


Figure 12.6 ❚ Because solids do not have variable concentrations, the equilibrium constant for this reaction is just the concentration of CO2. This means that the CO2 pressure at a given temperature is independent of the amounts of CaCO3 and CaO present. Here, the CO2 pressure is the same over two different mixtures of the two solids.

Same pressure of CO2(g)

CaCO3(s) CaO(s)

Much CaCO3(s), little CaO(s)


Little CaCO3(s), much CaO(s)

12 . 2

Calcium hydroxide will precipitate from solution by the following equilibrium: Ca21(aq) 1 2 OH2(aq) N Ca(OH)2(s) Write the equilibrium expression for this reaction.

Strategy Write the equilibrium expression as before but do not include a term for the concentration of the solid calcium hydroxide product.

Solution 1 K5— 21 Ca OH2 2 c

dc

d

Analyze Your Answer This answer may seem troublesome because there are no concentration terms in the numerator. But when we look at the reaction, we see that the only product is a solid. Because pure substances do not appear in the equilibrium constant expression, we are left with just a “1” in the numerator, and the answer is correct.

Check Your Understanding Write the chemical equation for the precipitation of copper(II) phosphate from a solution containing copper(II) and phosphate ions and then write the corresponding equilibrium expression.

Numerical Importance of the Equilibrium Expression Being able to write an expression relating equilibrium concentrations in chemical reactions provides a powerful tool for understanding chemical equilibrium. We will soon see how to evaluate these expressions for several classes of reactions. But first, let’s look at what we can learn from the numerical values of equilibrium constants. For example, we could ask the question, “Which calcium compound is more likely to form if more than one anion is present?” How can we use equilibrium constants to answer this type of question? Equilibrium constants range over many orders of magnitude. Values as small as 10299 are known, as are values as high as 1099 or greater. What does the size of the equilibrium constant tell us about the reaction? If we look at the equilibrium expression for a reaction as an algebraic relationship, the answer will become apparent. [products] K5— [reactants]

Sometimes, the power of 10 in an equilibrium constant exceeds 99, causing errors on some calculators. Such cases can be handled by algebraically manipulating the exponents.

12.3


401

Concentrations of products are in the numerator, and concentrations of reactants are in the denominator. For a very large value of K, the product concentrations must be large and the reactant concentrations must be small. By contrast, for the value of K to be small, concentrations of the reactants must be large and those of products must be small.

[products]

K 5 — 5 large value [reactants]

[products] K 5 — 5 small value

[reactants]

Thus, the value of K can tell us the direction in which a chemical reaction is likely to proceed. For large values of K, products are favored, whereas small values of K (meaning K 90% 235U, whereas that used in nuclear reactors contains 3–5% 235U. Pellets of enriched uranium oxide are embedded in fuel rods and covered with water (Figure 14.9) in the reactor core. To initiate fission, a source of neutrons must be incorporated into the reactor design. Once this neutron source initiates fission, the chain reaction is self-sustaining. Control rods composed of cadmium or boron absorb extra neutrons, regulating the number of neutrons that impact 235U nuclei and maintaining a steady rate of fission. Control rods can be inserted between the fuel rods to slow or stop the reaction. The water surrounding the fuel rods cools and moderates the

492

Chapter 14 Nuclear Chemistry

Control rods absorb neutrons and regulate the rate of fission.

Cooling tower Heat exchanger

shown here is used in all U.S. nuclear power plants.

Condenser (cold water condenses steam from turbine)

Steam Control rod

Figure 14.9 ❚ The general design

Pump

Nuclear reactor Warm water

Uranium fuel rod

Cold water Molten sodium or liquid water under high pressure carries heat to heat exchanger.

Steam turbine generates electricity.

reactor; fast neutrons interact with water molecules and slow down, so that effective collisions can occur with 235U nuclei. Other materials, including graphite and “heavy water” (deuterium oxide, D2O, also written 2H2O), are sometimes used as moderators in different types of reactors. The cooling water carries the heat released in the reaction to the steam turbine. Steam turns the turbine, and electricity is generated.

Nuclear Waste Several of the products of fission in a nuclear reactor are themselves radioactive isotopes. Much of this radioactive material is concentrated in used or “spent” fuel rods. Because of the large amounts of radioactivity in these rods, they are referred to as high-level nuclear waste. Some of the radioisotopes involved have very long half-lives, so the storage or disposal of this high-level waste presents enormous challenges for both engineering and public policy. It is possible to reprocess spent rods to produce new fuel pellets. But this process has not been carried out in the United States since the 1970s because of regulatory concerns and nuclear nonproliferation treaties. All high-level waste produced in U.S. nuclear reactors is currently stored on-site at the reactor. But clearly this practice cannot continue indefinitely and raises its own safety and security concerns. In 1982, Congress established the Nuclear Waste Policy Act mandating that the U.S. Department of Energy find and develop an appropriate underground disposal facility. In 1987, Yucca Mountain in southwest Nevada was chosen as the site for an extensive feasibility study (Figure 14.10). Yucca Mountain has several characteristics that make it a possible site for burying high-level waste. It is extremely remote, the climate is dry, and the water table is about 1000 feet below the potential burial vault. Site characterization studies continue at Yucca Mountain, but many opponents of this site actively protest against the development of the facility. The U.S. Congress has not been able to muster the votes needed to proceed with the development of the Yucca Mountain site, and the long-term storage of high-level waste remains a contentious issue. Beyond the political and policy aspects of this issue are also tremendous engineering challenges. An acceptable storage facility for high-level nuclear waste will have to be designed so that it can remain intact for thousands of years. Even predicting environmental conditions over such a timescale is obviously highly speculative. Moreover, the construction materials will need to withstand the effects of high levels of radiation exposure over their long lifetime. As we’ll see in Section 14.7, radiation can have significant effects on many materials.

14.6 Transmutation, Fission, and Fusion

493

Figure 14.10 ❚ The proposed Yucca Mountain storage facility for high-level nuclear waste is shown in this artist’s drawing. Waste would be stored in a three square mile network of tunnels 1000 feet below ground level.

Hoist building

Shaft for people, material

Filter house

Nuclear waste receiving building

Administration building

Waste shaft Ventilation shaft

1000 feet Tunnel network Mine level

Tunnel

Rock strata

Steel waste canisters containing spent nuclear fuel

Fusion Nuclear fission has been harnessed to produce electricity for decades, but nuclear fusion has yet to realize its potential as a relatively safe source of energy. In a fusion reaction, small nuclei combine to form larger, more stable nuclei. The energy of the sun originates in a fusion reaction in which four hydrogen nuclei combine to form a helium nucleus. 4 11 H : 42 He 1 2 01 b 1 2n 1 energy However, as we saw in our discussion of chemical kinetics (Section 11.6), the simultaneous collision of four particles is very unlikely. Instead, the reaction proceeds in a stepwise fashion, one hydrogen nucleus colliding with another to produce an intermediate that collides with a third proton and so on. The reaction of four protons is too slow to be used in a fusion reactor. A more practical approach is to fuse two heavier isotopes of hydrogen: deuterium (2H) and tritium (3H): 2 1H

1 31 H : 42 He 1 10 n

This reaction produces more energy per nucleus than fission, and it is attractive for a number of reasons. Deuterium is a naturally occurring isotope; it makes up 0.015% of all hydrogen atoms. So the available supply of deuterium is practically unlimited. Tritium can be produced from 6Li: 6 3 Li

494


1 10 n : 42 He 1 31 H

Fusion does not produce the high-level radioactive waste that fission generates. The primary product is ordinary helium, which poses no radiation threat. The fact that energetic neutrons are also produced does pose some problems. Neutron bombardment could induce nuclear reactions in materials surrounding the fusion reactor, producing some level of radioactivity. But this also exists in fission reactors, and careful choices of engineering materials can minimize the risks involved. The use of fusion to generate electricity is complicated by a number of factors, one of which is the repulsion of positively charged nuclei. Enormous energy is needed to force nuclei close enough together to overcome the coulombic forces between protons, so the reaction can be initiated only at temperatures on the order of 106 K. In the hydrogen bomb—the first application of fusion—the initiator was a fission bomb that forced the nuclei together. In controlled fusion, the means of initiating the reaction must not be destructive. To produce an economically viable energy source, it must require less energy input than will be released by fusion. Otherwise, energy is consumed by the process rather than produced. Finally, with temperatures reaching 106 K, the reaction has to be confined somehow despite the fact that any solid material would melt. Taken together, these factors present enormous engineering challenges. Two means of solving the confinement problem inherent in controlled fusion are showing promise. One is magnetic confinement in which the high-energy plasma produced at very high temperatures is controlled in a magnetic field. Another is inertial confinement in which a pellet of fuel is dropped into a chamber and imploded by high-energy lasers. So far, although energy has been produced in fusion reactors, getting the reaction to go forward has required more energy than it produced. Fusion research continues, and an international collaboration is finalizing plans to build the next generation reactor, known as ITER, in Cadarache, France. (The name ITER comes from the Latin word iter, which means “the way.”) It is estimated that fusion may become a viable source of energy late in the twenty-first century.

14.7

In kinetic terms, fusion has an extremely high activation energy.

The Interaction of Radiation and Matter

The effects of radiation on matter are governed by three main factors. The first is simply the amount of radiation to which matter is exposed. As you would probably expect, higher doses of radiation usually have more serious effects. But the impact of any radiation exposure also depends critically on two characteristics of the radiation itself, commonly referred to as penetrating power and ionizing power.

Ionizing and Penetrating Power of Radiation Any radiation can be classified as either ionizing or nonionizing; the distinction is based on the energy carried by a photon or particle. If that energy is greater than the ionization energy of typical atoms or molecules, then we should expect that the radiation could induce ionization in whatever material it encounters. Nonionizing radiation includes visible light, radio waves, and microwaves; all have photon energies smaller than typical ionization energies. X-rays and gamma rays, on the other hand, have much larger photon energies, and so they are ionizing radiation. Alpha and beta particles are also ionizing radiation. Ionizing radiation is much more likely to cause significant damage to any materials that it encounters, including living tissue. When ionizing radiation passes through living tissue, it can eject electrons from atoms or molecules it encounters, producing free radicals (Section 2.8). These highly reactive particles, either atoms or molecular fragments, can then scavenge electrons from other molecules. When free radicals react in living tissue, results can range from surface burns to serious genetic damage, as might occur if part of a DNA molecule 14.7 The Interaction of Radiation and Matter

495

is changed. One commonly formed molecular fragment is the extremely reactive hydroxyl radical, O9H. Once formed in a cell, the hydroxyl radical will react with nearby molecules, and the cell may be irreparably damaged. Thus, ionizing radiation can lead directly to cell death. Ionizing power alone cannot predict the impact of radiation on matter, however. We must also consider penetrating power—how far a particle penetrates into a material before its energy is absorbed or dissipated. The high energies of alpha particles, for example, give them large ionizing power. But because an alpha particle is relatively large, it does not tend to penetrate deeply into matter. Alpha particles can be stopped by cells in the outer layer of skin, where they pick up electrons to become helium atoms. Because the energy of the alpha particle is dissipated primarily in the skin, it can cause surface burns. But alpha particles originating outside of the body usually do not cause more serious harm because they do not penetrate sufficiently to reach internal organs. Alpha particles produced inside the body are a much greater danger because their energy will be deposited in internal organs. This is the primary risk from radon gas. Radon is a naturally occurring gas, and the amount in any area varies based on the types of soil and rocks present. As a member of Group 18—the noble gases—radon is chemically inert. But it is radioactive and decays by alpha emission.

Figure 14.11 ❚ The possible health hazards from exposure to ionizing radiation depend on the penetrating power of the radiation. Alpha particles are stopped in outer layers of the skin, whereas beta particles penetrate several centimeters into the body. Gamma rays can pass completely through the body.

496

: 42 He 1 218 84 Po

218 84 Po

: 42 He 1 214 82 Pb

Because radon is a gas, it can be inhaled. If an inhaled radon atom undergoes decay, alpha particles will be produced in the lungs. The dissipation of the energy of those alpha particles can cause serious tissue damage, leading to increased risk of lung cancer. Beta particles typically have lower energy than alpha particles, and so they might seem safer. But because beta particles are smaller, they can pass several centimeters into the body. Because of this greater penetrating power, beta radiation is often more dangerous than alpha radiation. Gamma rays can pass entirely through the body, and their interactions with atoms and molecules in vital organs can do a great deal of damage. Figure 14.11 illustrates the penetrating power of various forms of ionizing radiation. As our discussion thus far implies, ideas such as ionizing and penetrating power are important in assessing the possible health effects of radiation exposure. But these same ideas also govern the effect of radiation on materials or electronic devices and can have important engineering implications. In the opening section of this chapter, we learned that the distribution of cosmic rays entering the atmosphere is very different from that reaching Earth’s surface. So electronic devices in space are exposed to a much higher level of radiation than they would be at ground level. As modern communications have become more dependent on satellite-based electronics, the need to study and protect against the possible effects of cosmic rays has grown. Computer chips and other solid-state electronic devices rely on a carefully controlled distribution of electrons and holes in semiconductor materials. So the

X-rays and gamma rays penetrate deep into tissue. Because they can pass completely through the body, they can damage internal organs.

X-rays and gamma rays

Beta particles penetrate into the skin but generally cannot reach internal organs.

Beta particles

Alpha particles have very low penetrating power, and are stopped by the outer layers of skin.

Alpha particles


Adapted from Health Care Protection Office, The University of Iowa

See Section 8.3 for an introduction to semiconductors.

222 86 Rn

Cathode

Thin window Radioactive sample (e.g., uranium ore)

Anode

Argon gas

Voltage source

Amplifier

Speaker

Geiger counter hand piece

Figure 14.12 ❚ In a Geiger-Mueller tube, or Geiger counter, radiation passes through a thin window into a gas-filled tube. Energy from the radiation produces ions in the gas, releasing electrons. The resulting ions and electrons are then attracted to oppositely charged electrodes, producing a pulse of electric current. In a typical Geiger counter, the current pulse is converted to an audible clicking sound.

production of ions within those materials can cause catastrophic failures. The cosmic rays most likely to induce ionization are heavy nuclei at high energy, which are much more likely to be encountered in space than at ground level. Even a single such particle can produce a large number of ions, leading to what is known as a single event effect. In some cases, these effects can be temporary and perhaps require resetting a device. But other single event effects can be much more serious, causing permanent hardware damage. As a result, chips for use in satellites are frequently packaged in special hardened materials, designed to protect them from the cosmic rays they will encounter. Dangers from exposure to cosmic rays also represent a significant obstacle for long-term manned space missions, such as travel to other planets.

To assess radiation doses, we must have suitable methods for measuring the type and amount of the radiation. Rutherford’s early experiments used a fluorescent screen coated with a zinc sulfide (ZnS) phosphor. The fluorescent screen produced a pinpoint of light each time it was struck by an alpha particle. The tiny flashes of light had to be observed and counted by Rutherford and his assistants, an arduous and potentially inaccurate task. Such a method is clearly not suitable for modern applications in medical imaging or carbon dating. The scintillation counter is a modern adaptation of the visually recorded light pulse on a screen. Such a counter uses a fluorescent material to detect radiation, which produces photons. Those photons can then be counted using a photomultiplier tube and appropriate electronics. The portable Geiger counter is commonly used in laboratories and by radiation safety teams to measure radioactivity (Figure 14.12). A glass tube containing a gas at a low pressure (about 0.1 atm) is coated on the inside with a metal that acts as a cathode. A wire anode runs down the center of the tube. A high voltage is applied across the electrodes. Alpha and beta particles enter a window in the tube and ionize atoms of gas. Electrons released from the gas atoms are attracted to the anode, and, as they travel to the anode, they cause the ionization of more gas atoms, releasing more electrons. An avalanche of electrons reaches the anode, and a current pulse is recorded. If it is connected to a speaker, the pulse is registered as an audible click. A film-badge dosimeter, commonly worn by people who work with radioactive isotopes, takes advantage of Becquerel’s discovery that radiation darkens photographic plates (Figure 14.13). The darkened badge, coupled with a record of each exposure to a radioactive species, provides a warning mechanism if safe exposure levels are exceeded.

Lawrence S. Brown

Methods of Detecting Radiation

Figure 14.13 ❚ People who work around radioactive materials wear film badges (like the one shown here) to monitor their exposure to radiation.

14.7 The Interaction of Radiation and Matter

497

Table

❚

14.2

Definitions and units used to quantify exposure to radiation

US

SI

Exposure—an older unit for measuring exposure to photons (gamma and X-rays) in dry air. It is a measure of the ionization of molecules of air in coulombs/kg of air.

Roentgen (R) 1 R 5 2.58 3 104 C/kg of dry air

Absorbed Dose—amount of any type of radiation absorbed in any kind of material. Doesn’t describe biological effects of different types of radiation.

Radiation absorbed dose (rad) 1 rad 5 1 erg/g (1 erg 5 1022 J)

Gray (Gy) 1 Gy 5 100 rads

Equivalent Dose—relates the absorbed dose to the damage produced by different types of radiation. A quality factor, Q, is multiplied by the absorbed dose to get the equivalent dose.

Roentgen equivalent man (rem) 1 rem 5 Q 3 absorbed dose in rad

Sievert (Sv) 1 Sv 5 100 rem 1 Sv 5 Q 3 absorbed dose in Gy

Methods of detecting radiation must properly account for background radiation. Cosmic rays are a constant source of radioactivity, for example, as are natural radioactive isotopes in soil, air, and water. This radiation is always around us and must be subtracted from measurements of radioactive sources.

Measuring Radiation Dose Radiation detectors such as those described above give us the ability to quantify the level of radiation in a particular environment. Because of the interplay between ionizing power and penetrating power, there are a number of different ways to express radiation dose. The exposure measures the number of ions produced in air. The absorbed dose, as the name implies, measures the amount of radiation actually absorbed by a particular material. The equivalent dose attempts to quantify the resulting damage to human tissue. To further complicate things, two different sets of units are commonly used to account for the differences in ionizing power and penetrating power in assessing the extent of effective damage in human tissue. These quantities and their respective units are summarized in Table 14.2. The quality factor, Q, used in calculating the equivalent dose is also known as the relative biological effectiveness (RBE). It varies from a value of one for high-energy photons (gamma and X-rays) to about 20 for alpha particles. The definition of equivalent dose attempts to incorporate information about exposure and toxicity simultaneously.

INSIGHT INTO

14.8 Radiation is also used to treat various cancers.

498

Modern Medical Imaging Methods

Radiation has been used as a tool for medical diagnosis since the introduction of X-rays as a way to see into the body. More recent developments include the use of various radioactive isotopes to obtain images of specific organs, as well as more elaborate techniques such as positron emission tomography (PET). In a traditional X-ray, radiation passes through the body and a photographic image is produced based on the extent to which the radiation is absorbed. Because


Courtesy of The National Institute on Aging

bones absorb X-rays much more strongly than organs or other tissues, X-ray images are excellent tools for orthopedic diagnosis. X-ray images can also be used to examine the structure of some organs, as in a chest X-ray to examine the heart or lungs. But they are not generally suitable for studying the function of organs. Instead, a range of techniques has been developed around the idea that if a small amount of an appropriate radioisotope can be selectively deposited in the target organ, then monitoring the radiation from that isotope can produce a detailed image of the organ. If such images can be taken with sufficient time resolution, it becomes possible to monitor processes such as blood flow or oxygen uptake, providing a powerful and relatively noninvasive medical tool. How can radioactive isotopes be introduced into specific organs? The answer relies on the fact that biochemistry tells us that certain elements and compounds are taken up specifically by particular organs. The thyroid gland uses iodine to produce thyroid hormone, for example. So, radioactive 131I is often used to diagnose thyroid problems. The patient is given an injection of a small amount of the radioisotope, and then natural biochemical pathways carry the 131I to the thyroid. 131I undergoes beta decay, so detection of radiation emanating from the thyroid can be used to produce an image of the gland. Here the organ itself becomes the light source, so only an external radiation detection scheme is needed to record the image. Several other isotopes are used in similar ways to image different organs. So long as the radiation dose is fairly small and the half-life of the isotope being used is not too long, these procedures are extremely safe. As the name implies, PET images (Figure 14.14) are based on isotopes that emit positrons. Recall from Section 14.4 that positron emission tends to occur in neutrondeficient isotopes. Among the available positron emitters are 11C, 18F, 13N, and 15O. Because all of these elements are found in common organic molecules, it is relatively easy to incorporate them into appropriate biological molecules. 15O might be incorporated into glucose, for example. Because our bodies rely on glucose as a source of energy, it will be taken up by most organs. Each decay of the radioisotope releases a positron, but the lifetime of a positron in the body is extremely short. In most cases, a positron will travel no more than a couple of millimeters before it encounters an electron. The positron and electron then undergo matter2antimatter annihilation, which releases a pair of gamma-ray photons traveling in opposite directions. Detectors register the gamma rays, and computers map out the path taken by the tagged compound. The result is a map of a slice through the body. The half-life of 15O is just 122 s, so decay to stable (and harmless) 15N is quick. This short half-life is beneficial because the radiation is only in the body for a very short time. But it also provides some challenges in carrying out the procedure because the short-lived radioisotope cannot be stored. So, PET scans can be performed only in facilities that can produce the needed isotopes on-site.

Figure 14.14 ❚ Positron emission tomography (PET) produces high-quality images of the brain and other organs. The technique is based on the simultaneous detection of pairs of gamma ray photons produced when a positron and an electron collide. This makes it possible to measure even low signal levels very accurately.

14.8 Modern Medical Imaging Methods

499

FOCU S O N

PRO BLEM S O LVI NG

Question Suppose that you are working for a company that prepares radioisotopes

for medical imaging. You receive an order from a researcher requesting 15 mg of 45Ti for an experiment. The customer’s laboratory is located roughly 1.5 hours from your production source. How would you determine what mass of the isotope to produce to fill this order?

Strategy The key issue here is that the isotope you produce will decay during the time that it is in transit to the research lab. Because the experiment requires a set mass of the isotope, you must consider the kinetics and produce more than is required so that there will be enough 45Ti left upon delivery. The most important piece of information to look up is the half-life of 45Ti. Once that value is known, we can use the integrated rate law for first-order kinetics to carry out the needed calculations.

Solution Once we know the half-life, we can determine the rate constant, k: 0.693 k 5 ——— t 1/2

Next we turn to the integrated rate law for a first-order reaction. In Equation 14.1, we wrote this as N 5 N0e2kt Here we can use masses in place of the N and N0 terms, giving us m 5 m0e2kt Our goal is to find m0, the initial mass of 45Ti needed. We have to make an assumption about the delivery time and the amount of time it would take to package the isotope for delivery. If we assume 30 minutes from production to getting the isotope aboard the delivery vehicle, we might use 120 minutes as our time. m 5 m0e2kt 15 mg m ————— m0 5 —— 5 e2kt e2k(120 min)

S U MMA RY Nuclear reactions involve the conversion of one chemical element into another and are typically accompanied by the emission of radiation in the form of high-energy particles and photons. Some nuclear reactions occur naturally, as in radioactive decay, whereas others can be induced by human intervention. Nuclear decay processes can be categorized in terms of the type of radiation emitted. We can write equations for nuclear reactions similar to those for ordinary chemical reactions. But in a nuclear reaction, the number of atoms of each element is not conserved, so we must use different ideas to balance the equations. In a properly balanced nuclear equation, both mass number and charge number should be conserved. All radioactive decay processes follow first-order kinetics, and their rates are commonly described in terms of their halflives. The standard equations for first-order reactions can thus be used to predict the activity of a radioactive sample as a function

500


of time. Rates of decay are usually reported in units of becquerels or curies. Carbon dating relies on the fact that the kinetics of the decay of 14C are well understood. Nuclear chemists rely on the chart of the nuclides to summarize their knowledge of the stability of various isotopes. All known stable isotopes fall into a region known as the band of stability, and the position of unstable isotopes with respect to this band can be used to predict the type of decay expected. Compared to ordinary chemical reactions, nuclear reactions release tremendous amounts of energy. The energy change of a nuclear process can be calculated from the mass defect by using the Einstein equation, E 5 mc2. The same type of calculation allows us to find the binding energy for any given nucleus and reveals that nuclear binding energy goes through a maximum for 56 Fe. Thus very heavy nuclei can undergo fission, splitting into two or more lighter nuclei with increased stability. Lighter nuclei, in contrast, may gain stability by undergoing fusion, in which

two or more nuclei combine to form a larger and more stable nucleus. Current nuclear reactor technology is based on the fission of 235U. Fusion reactors also hold great promise as future energy sources but still face significant engineering challenges.

The effects of radiation on matter—including living tissues—depend on both the type of radiation and the extent of exposure. In assessing its effect on matter, radiation is characterized in terms of ionizing power and penetrating power.

K EY TE RM S absorbed dose (14.7)

critical mass (14.6)

mass defect (14.5)

alpha decay (14.2)

curie (Ci) (14.3)

nuclear fission (14.5)

alpha particle (14.2)

decay constant (14.3)

nuclear fusion (14.5)

alpha ray (14.2)

decay series (14.4)

nuclear reaction (14.1)

background radiation (14.7)

dosimeter (14.7)

nuclide (14.2)

band of stability (14.4)

electron capture (14.2)

penetrating power (14.7)

becquerel (Bq) (14.3)

electron volt (eV) (14.1)

positron (14.2)

beta decay (14.2)

equivalent dose (14.7)

positron emission tomography (PET) (14.8)

beta particle (14.2)

exposure (14.7)

quality factor (14.7)

beta ray (14.2)

fuel rod (14.6)

radioactive decay (14.1)

binding energy (14.5)

gamma ray (14.2)

reactor core (14.6)

chain reaction (14.6)

Geiger counter (14.7)

relative biological effectiveness (RBE) (14.7)

chart of the nuclides (14.4)

half-life (14.3)

sea of instability (14.4)

compound nucleus (14.6)

ionizing power (14.7)

strong force (14.4)

control rod (14.6)

K capture (14.2)

transmutation (14.6)

cosmic ray (14.1)

magic number (14.5)



INSIGHT INTO Cosmic Rays and Carbon Dating 14.1 Cosmic rays are sometimes referred to as corpuscular rays. How does this term distinguish them from other rays of the sun that reach the planet? 14.2 Who would have greater exposure to cosmic rays: a commercial pilot or an engineer who helps build the runways the pilot uses? Explain your answer.

Radioactivity and Nuclear Reactions 14.7 Refer to Figure 14.1 and describe the experiment in which Ernest Rutherford demonstrated that different types of radiation emanate from uranium. 14.8 List two ways in which a gamma ray differs from an alpha or beta particle. 14.9 Match the following forms of radioactive decay with the appropriate result.

14.3 Use the web to find information on the composition of cosmic rays beyond hydrogen and helium.

(a) Alpha

1. No change in mass number or atomic number

14.4 A fastball thrown by a professional baseball pitcher travels at about 100 mph. How does the kinetic energy of the baseball thrown at this speed compare to cosmic radiation at 15 MeV? (The mass of a baseball is about 145 g.)

(b) Positron

2. Atomic number decreases by 1

(c) Gamma

3. Atomic number decreases by 2, mass number decreases by 4

(d) K capture

4. Atomic number increases by 1

14.5 (a) How does carbon-14 enter a living plant?

(e) Beta

(b) Write the equation for this reaction. 14.6 Explain the following: (a) The ratio of 14C to 12C in a living plant is constant. (b) The constant ratio cited in (a) is a steady-state condition, not an equilibrium.

■


14.10 Why does positron decay have the same effect on the nucleus as electron capture? 14.11 Complete each equation and name the particle ejected from the nucleus. __ (a) 208 O : ? 1 210 b 1 n


501

14.12

(b)

?

(c)

201 82 Pb

4 : 232 92 U 1 2 He

:

201 83 Bi

Complete the following nuclear equations. Write the mass number, atomic number, and symbol for the remaining particle.

(b) Compute the activity of a 1.000-g sample of 110Sb in Bq and Ci.

■

4 (b) ?1 10 n : 24 11 Na 1 2 He 40 1 (c) 40 20 Ca 1 ? : 19 K 1 1 H 243 4 (d) 241 95 Am 1 2 He : 97 Bk 1 ? 12 1 (e) 246 96 Cm 1 6 C : 4 0 n 1 ? 249 1 (f) 238 92 U 1 ? : 100 Fm 15 0 n

14.13 Write equations for the following nuclear reactions: (a) Alpha decay by 188Bi (b) Beta emission by 87Rb (c) Positron emission by 40K (d) Electron capture by 138La 14.14 Write equations for the following nuclear reactions: (a) 94Tc undergoes electron capture.

14.22

■ Gold-198 is used in the diagnosis of liver problems. The half-life of 198Au is 2.69 days. If you begin with 2.8 mg of this gold isotope, what mass remains after 10.8 days?

14.23 The half-life of 19O is 29 s. Suppose that a scientist wishes to incorporate 19O into a molecule and then use its radioactivity to trace the fate of that molecule in a reaction. To make an accurate measurement, there must be at least 1.50 mg of 19O left by the end of the study, which will take 2.5 min each time the experiment is run. What is the minimum mass of 19O that must be used each time the experiment is conducted? 14.24 Archaeologists use the year 1950 as the date of the present. So 1000 b.c. is 2950 b.p. (Before the Present). An archaeologist finds that the carbon-14 content of a wooden needle is 11% of that of trees living today. At what date b.p. was the tree from which the needle was made cut down?

(c) 217Fr emits an alpha particle.

14.25 A wooden artifact is burned and found to contain 21 g of carbon. The 14C activity of the sample is 105 disintegrations/min. What is the age of the artifact?

(d) 220Ra is formed by alpha emission.

14.26

(b) 89Nb ejects a positron.

14.15 (a) Explain why tritium, 3H, cannot undergo alpha decay. (b) What type of decay would you expect tritium to undergo? (c) Write the nuclear equation for the expected decay. ■

By what processes do these transformations occur?

(a) Thorium-230 to radium-226 (b) Cesium-137 to barium-137 (c) Potassium-38 to argon-38 (d) Zirconium-97 to niobium-97 14.17 One way to convert lead into gold is to irradiate 206Pb with neutrons. There are several steps in the process, and three neutrons are required for each lead atom. In addition to gold, three beta particles and three alpha particles are produced. (This process requires a particle accelerator, and the cost is much higher than the value of any gold that could be produced.) (a) What isotope of gold will this process produce? (b) Write the overall nuclear reaction for the process. Radioactive Decay Rates 14.18

■ What is the half-life of a radioisotope if it decays to 12.5% of its radioactivity in 12 years?

14.19

137

Cs has a half-life of 30.2 years. How many years will it take for a 100.0-g sample to decay to 0.01 g?

14.20

■

Cobalt-60 is used extensively in medicine as a source of g-rays. Its half-life is 5.27 years.

(a) How long will it take a Co-60 source to decrease to 18% of its original activity? (b) What percent of its activity remains after 29 months?

502

(a) Determine its decay constant in s21.

1?

(a) 94 Be 1 ? : 63 Li 1 42 He

14.16

14.21 The half-life of Sb-110 is 23.0 s.


■ Some mammoth bones found in Arizona were found by carbon-14 dating to be 1.13 3 104 years old. What must have been the activity of the carbon-14 in the bones in disintegrations per minute per gram? Assume the original activity was 15.3 disintegrations per minute per gram.

14.27 Tritium, 3H, is a radioactive isotope of hydrogen produced by cosmic rays in the atmosphere, where it oxidizes and becomes part of the water cycle. Its half-life is 12.33 yr. Which of the following could be dated using tritium? Explain your answer. (a) 4000-year-old ice from Antarctica (b) Modern French wine (c) Ancient Egyptian beer from a tomb in the Valley of the Kings 14.28 Explain why uranium can be used for geological dating but not for dating artifacts from human activity on Earth. 14.29 A piece of a spear handle is found in an archaeological dig in Central America. It contains 12.5% as much 14C as a tree living today. Based only on the half-life of carbon-14, how old is the spear handle? 14.30 Beginning in the late nineteenth century and continuing into the twentieth century, the Industrial Revolution was fueled by coal and spewed enormous volumes of 14Cdepleted CO2 into the atmosphere. (a) Why is coal 14C-depleted? (b) How would the increase in 14C-depleted CO2 affect the date of a very old object in which its 14C/12C ratio was compared to the 20th-century ratio? Explain. Nuclear Stability 14.31 What is the N/Z ratio for each of the following nuclides? (a) 147 N

(b) 114 50 Sn

(c) 234 90 Th

■


14.32

Write a nuclear equation for the type of decay each of these unstable isotopes is most likely to undergo.

14.41 Compute the binding energy of the 7Li nucleus, whose experimentally determined mass is 7.016004 u.

(a) Neon-19

14.42 Find the mass defect in the reaction in which a neutron decays to a proton and an electron.

■

(b) Thorium-230

14.43 Compute the binding energy of the 14N nucleus, which has an experimentally determined mass of 14.003074 u.

(c) Bromine-82 (d) Polonium-212 14.33 Based on their positions relative to the band of stability, predict the type of decay that each of the following will undergo. Write equations for the expected decays. (a) 73 36 Kr

(b) 138 O

(c) 126 50 Sn

14.34 The following successive decays occur, starting with 24 F, which has a half-life of 0.3 s, and ending at a stable isotope: 24 9F

,t1/2 0.3 s

?

,t1/2 3.4 m

?

,t1/2 15 h

?

(a) Identify each of the decay products. (b) Using the chart of the nuclides, explain why successive beta emissions make more sense than (i) alpha emission; (ii) positron emission; or (iii) electron capture in producing a stable nucleus from 24F. 14.35

The thorium-232 radioactive decay series, beginning with 232Th and ending with 208Pb, occurs in the following sequence: a, b2, b2, a, a, a, a, b2, b2, a. Write an equation for each step in this series.

■

14.36 Starting with 238U and ending with an isotope of Pb, write a decay scheme in which every nuclide emits only an a particle. Compare your scheme to the decay series in Figure 14.5 and suggest a reason why b2 emissions must occur in the 238U decay series. 14.37 The 238U series can be used to date minerals containing uranium by measuring the ratio of 238U to 206Pb, but all of the decay products have to remain in the mineral through several of their half-lives so that essentially all the 238U that decays becomes 206Pb. One of the decay products in this series is the noble gas, 222Rn. How would the experimentally determined age of the mineral be affected if some of the radon diffused from the rock? 238

232

222

14.38 Both U and Th decay to stable isotopes of lead. Rn (t1/2 5 3.8 d) is an intermediate nuclide in the 238U series. 218 Rn (t1/2 5 56 s) is an intermediate in the 232Th decay series. In which series would diffusion of gaseous radon cause the greater error in the computed age of a rock (see Exercise 14.37)? Energetics of Nuclear Reactions 14.39 Calculate the binding energy of 1 mole of 14C nuclei. The experimentally determined mass of a carbon-14 atom (including its six electrons) is 14.003242 u. 14.40

Collision of an electron and a positron results in formation of two g-rays. In the process, their masses are converted completely into energy.

■

(a) Calculate the energy evolved from the annihilation of an electron and a positron in kilojoules per mole. (b) Using Planck’s equation (Equation 6.2), determine the frequency of the g-rays emitted in this process.

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14.44

The actual mass of a 108Pd atom is 107.90389 amu. (a) Calculate the mass defect in amu per atom and in g per mole for this isotope. (b) What is the nuclear binding energy in kilojoules per mole for this isotope?

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14.45 It takes 360 kJ to keep a 100-watt lightbulb burning for 1 hour. Assuming that the lightbulb won’t burn out, for how many years could the energy of the mass defect of one mole of 14C keep a 100-W lightbulb burning? 14.46 It has been suggested that element 114 has unusual stability, but 114 is not a magic number. What is a possible stable neutron number for this element? 14.47 Lead has a magic number of protons. A doubly magic nuclide should be exceptionally stable. Why isn’t lead-164 stable? 14.48 Tin has a magic number of protons. One of its isotopes, 132 Sn, has a magic number of neutrons, as well, but it is unstable. Explain. 14.49 Use the web to identify the heaviest element that has been observed to form a chemical compound. What element is it and what compound does it form? Are heavier elements incapable of forming compounds? 14.50 Glenn Seaborg was a leader in the search for the heaviest elements, and element 106 is named Seaborgium (Sg) in his honor. In an article in the Journal of Chemical Education, Seaborg summarized the history of the search for the transuranium elements up to the mid-1980s (Glenn T. Seaborg, “The Transuranium Elements,” Journal of Chemical Education, vol. 62, 1985, pp. 4632467). Read the article and answer the following questions. (a) Who was responsible for the position of the lanthanides in the periodic table? (b) What did Enrico Fermi think was being produced when he bombarded uranium with neutrons? (c) How did Ida Noddack’s paper straighten out the confusion about the identity of the products of Fermi’s nuclear reaction? (d) Compare Figures 1 and 2, the 1930 periodic table and Seaborg’s 1945 version of the table. What was Seaborg’s insight that gave us the modern periodic table? (e) State the two possible reasons Seaborg gives for the fact that the “island of stability” (element 114) has not been reached. (f) Why does fusion of two elements with heavy nuclei produce only a few atoms of a new, superheavy nucleus? (g) In answer to a student’s question, Seaborg announced to the world the discovery of two new elements in 1945 on the “The Quiz Kids,” a radio program in Chicago. Seaborg wrote, “This was the first time in the history of the world that the announcement of the discovery


503

of chemical elements was sponsored by Alka Seltzer.” What are the atomic numbers and names of the two elements? Transmutation, Fission, and Fusion 14.51 Ernest Rutherford carried out the first artificial transmutation in 1919 when he bombarded 14N with alpha particles. The result was an oxygen nucleus. 14 7N

1 42 He : 189 F* : 178 O 1 11 p

Radiochemists use a shorthand notation to represent transmutations. Omitting the intermediate 189 F* nucleus, the reaction above can be represented as 14N (a, p) 17O. Write balanced equations for each of the indicated nuclear bombardments. (a) 27Al (p, g) 28Si (b) 40Ar (n, b2) 41K 14.52 (a) Write the equation for the reaction that occurs when a 200Hg atom is bombarded by a proton. One of the products is an alpha particle. (Include the compound nucleus.) (b) Represent the reaction above in the shorthand notation introduced in the previous problem, omitting the compound nucleus. (c) The reaction above must take place in a particle accelerator so that the protons strike the mercury nuclei at very high energies. Why isn’t this method of producing the precious metal product in (a) the answer to the alchemist’s dreams? 14.53 Slow neutrons can be absorbed by some nuclei to produce new elements. Why does a neutron make a better projectile for a transmutation reaction than an alpha or beta particle? 14.54

What is the missing product in each of these fission equations?

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93 1 1 (a) 235 92 U 1 0 n : ?1 38 Sr 1 3 0 n 132 1 1 (b) 235 92 U 1 0 n : ?1 51 Sb 1 3 0 n 141 1 1 (c) 235 92 U 1 0 n : ?1 56 Ba 1 3 0 n

14.55 How much energy is released in the fission of 1 kg of 235 U according to the equation below? The experimentally determined masses of the products are 136.92532 u and 96.91095 u, respectively. (The masses of 235U and the neutron are given in Example Problem 14.6.) 235 92 U

97 1 1 10 n : 137 52 Te 1 40 Zr 12 0 n

14.56 Identify each of the following and indicate its purpose in a fission reactor. (a) moderator

(c) control rods

(b) fuel

(d) coolant

14.57 Because the fissionable nucleus of uranium is 235U, the large quantity of 238U in a nuclear reactor isn’t used. However, it is possible to produce fissionable 239Pu from 238 U by the reaction 238U (n, b2) 239Np. 239Np decays by 2 b emission to 239Pu. A reactor used for this purpose is called a breeder reactor. (a) What is the compound nucleus formed by 238U? (b) Write balanced equations for production of 239Pu.

504


14.58 (a) What mass of octane, C8H18, would have to undergo combustion to provide the same quantity of energy as that released by the fission of one kilogram of 235U to produce 90Sr and 143Xe? (b) How many gallons of octane is this? (The density of octane is 0.702 g/cm3.) 14.59 According to the U.S. Nuclear Regulatory Commission, in 2007 there were 104 commercial nuclear reactors in the United States. In 2007, 4.2 3 106 gigawatt-hours of electricity were generated for use, and 19.4% of that was generated by nuclear power. (a) A nuclear technician assumed that the products of the fission of 235U by one neutron were 97Zr and 137Te (see Problem 14.55). Ignoring decay products, the technician estimated the total mass of zirconium and tellurium in metric tons that would have to be disposed of in 2007 as nuclear waste (1 watt-hour 5 3600 joules; 1 metric ton 5 1000 kg). What was the technician’s result? (b) Using a chemical handbook, look up the half-lives of 97 Zr and 137Te. When the waste was disposed of, do you think there was a measurable quantity of either isotope? Explain. (c) In the handbook, look up the decay mode of each isotope and write the decay series to a stable isotope for each one. Which decay series has a radioactive isotope that will remain in the waste for many years? (d) Using the stable isotope products in your calculation, how far off was the technician’s estimate of total waste? Explain. 14.60 (a) Write the equation for the reaction of deuterium and tritium that could be used in a fusion reactor. (b) What are two obstacles to this reaction as a viable source of energy? 14.61 A fusion reaction that has been proposed involves bombarding boron-11 with a proton and producing a single, stable product. (a) Write the equation for the reaction. (b) Which of the obstacles in the previous problem is overcome by this reaction? 14.62 Use the Internet and other resources to find an estimate of the amount of electrical energy that may be required worldwide by the year 2020. Write a 2- to 3-page essay in which you briefly discuss the pros and cons of each of the following methods of electrical energy production to meet the estimated demand. (a) solar energy, (b) fossil fuels, (c) fission, (d) fusion 14.63 In 1998, the Nuclear Regulatory Commission suggested that residents who live near a nuclear power plant should have a supply of potassium iodide on hand in the event of a nuclear plant emergency. Explain why the suggestion makes sense. 14.64 The Yucca Mountain nuclear waste burial site is touted as a potential solution to the problem of storing highlevel nuclear waste. Using the Internet, find information about the site and, in a well-reasoned essay, formulate arguments for and against using Yucca Mountain to store nuclear waste. Based on your arguments, if building the site were your decision, what would you do?

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14.65 The Yucca Mountain site must last for roughly 10,000 years. What engineering uncertainties are introduced for the design of such a facility compared to a design for a tunnel for transportation services, for example? 14.66 What is the critical feature of nuclear fusion that requires special containment ideas, such as magnetic or inertial containment? The Interaction of Radiation and Matter 14.67 Use information from Section 6.7 to estimate which form of electromagnetic radiation is the lowest energy ionizing radiation. 14.68 Is it the production of ions that makes ionizing radiation dangerous to living tissue or some other effect? Explain your answer. 14.69 Rank the most common forms of radioactivity in order of increasing penetrating power. 14.70 Both penetrating power and ionizing power are important in determining the danger of radiation to living tissue. The goal of understanding damage to living tissues is to limit that damage. If you were an engineer working for a firm that designs safety clothes and apparatus for working in radioactive environments, (a) Under what conditions is ionizing power the more important factor to consider? (b) Under what conditions is penetrating power the more important factor? 14.71 Polonium-210 is an alpha-emitting nuclide found in cigarette smoke. How is the fact that it is found in such smoke related to its health hazard? 14.72 Why is ionizing cosmic radiation of greater concern for engineering designs on satellites than for ground-based electronics designs? 14.73 How is the damage from the possible impact of ionizing cosmic rays on an electronic device in a satellite related to the materials from which electronic devices are constructed? 14.74 A radioactive sample used in medicine registers an activity of 10 mCi. What is its activity in Bq? 14.75 A Geiger counter is set to click once for every 100 disintegrations that it detects. The counter registers 20 clicks per second for a sample of an unknown isotope. What is the activity of the isotope in Ci? In μCi? 21

14.76 You are a biologist studying Cu absorption from water in certain legumes. You have a small quantity of 67CuSO4 and a Geiger counter. ( 67Cu is a b2 emitter with a half-life of 61 h.) Design and describe an experiment that demonstrates whether the legume takes up Cu21 from water. 14.77 An average person is exposed to about 360 millirem of background radiation a year from a variety of natural and man-made sources. About two-thirds of that typically comes from inhalation of 222Rn produced in soil by the decay of 238U. Given that 222Rn decays by alpha emission, estimate the absorbed dose in (a) joules and (b) grays for a 60-kg person from one year’s inhalation of radon. 14.78 In the early part of the 20th century, watches were produced with hand-painted dials that glowed in the dark. The paint

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contained ZnS, phosphor, and radium, which decayed to produce alpha particles that excited the ZnS to glow when struck by a charged particle. The people (mostly women) who painted the dials often put a fine point on their paintbrushes by rolling the brush on their tongues. Many workers in this industry subsequently died of bone cancer and leukemia when their bone marrow was destroyed. Imagine that you are writing a letter from the future to a family member working in the early-20th-century watch industry and explain to your relative what the hazards are and what safety equipment your relative should use. 14.79 Concern over radiation exposure has led some local, state, and federal agencies to propose a zero dose limit above background radiation (see Problem 14.77). A zero dose limit means that industry could not allow anyone to be exposed to radiation above the average background exposure per year. Why does background radiation change with altitude? Suggest a reason why the airline industry might oppose a zero dose limit. 14.80 Plutonium-239 is produced in breeder reactors by successive reactions in which 238U is bombarded by neutrons and the product decays by two successive beta emissions. (a) Write the equations for the three nuclear reactions. (b) Inhalation of 239Pu, which decays by alpha emission, can cause lung cancer. Explain why inhalation of 239Pu may be more dangerous than getting an equal quantity of the isotope on the skin. INSIGHT INTO Modern Medical Imaging 14.81 Write the equation for the nuclear decay of each of the following isotopes used in positron emission tomography (PET): (a) 18F (b) 11C (c) 13N 14.82 In nuclear medicine, radioactive isotopes are often ingested or otherwise introduced into the human body. Discuss the properties you would expect medical radioisotopes to exhibit. 14.83 Write the nuclear equation for the antimatter annihilation of an electron by a positron to produce two gamma rays. 14.84 In boron neutron capture therapy (BNCT), 10B absorbs a neutron, producing a nucleus that decays by alpha emission. (a) Write the equations for these two successive events, including the compound nucleus. (b) Alpha emission by inhaled 239Pu causes lung cancer, but alpha emission in BNCT destroys cancerous tumors. Suggest a reason for this difference. 14.85 Osteoporosis (bone loss) is a growing medical problem as the average age of the U.S. population increases. It can be diagnosed using the isotope gadolinium-153. Using the Internet and other resources, research and write a paragraph about using 153Gd to determine bone density.

FOCUS ON PROBLEM SOLVING EXERCISES 14.86 A researcher wants to use radioactive metals to label an enzyme that is found in high concentration in cancer cells. Either 64Cu or 68Ga will bind with the enzyme. Suppose that you are the engineer assigned to this order by your company. What would you need to know about these


505

isotopes and what questions would you ask the researcher to help with the selection? 14.87 Thallium-201 has a long history of use in nuclear medicine. Suppose that a customer requires 140 mg of the isotope, but delivery will take 18 hours by an express courier service. You are assigned to determine how much isotope you need to produce. What information would you need to look up to answer the question? 14.88 A customer has ordered 70 mg of 18F radioisotope from your company. The delivery truck that normally would have made the trip from your factory to the customer in

506


15 minutes is delayed by a major traffic jam and will now take slightly more than an hour to arrive. If you packaged 100 mg of the isotope when it left the factory, will enough be left upon delivery? What information would you have to look up? 14.89 Suppose that you want to use radioactive iodine to run an experiment to treat a tumor. You can choose either 123I or 125 I for this work. What do you need to know about these isotopes? What design parameters for the experiment factor into your decision?

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Appendix A INTERNATIONAL TABLE OF ATOMIC WEIGHTS Atomic Number

Symbol

89

Ac

Actinium

[227]

72

13

Al

Aluminum

26.9815386

108

95

Am

Americium

[243]

2

He

Helium

4.002602

51

Sb

Antimony

121.760

67

Ho

Holmium

164.93032

18

Ar

Argon

39.948

1

H

Hydrogen

1.00794

33

As

Arsenic

74.92160

49

In

Indium

114.818

85

At

Astatine

[210]

53

I

Iodine

126.90447

56

Ba

Barium

137.327

77

Ir

Iridium

192.217

97

Bk

Berkelium

[247]

26

Fe

Iron

55.845

4

Be

Beryllium

9.012182

36

Kr

Krypton

83.798

Name

Atomic Weight

Atomic Number

Name

Atomic Weight

Hf

Hafnium

178.49

Hs

Hassium

[270]

Symbol

83

Bi

Bismuth

208.98040

57

La

Lanthanum

107

Bh

Bohrium

[272]

103

Lr

Lawrencium

138.90547 [262]

5

B

Boron

10.811

82

Pb

Lead

207.2

35

Br

Bromine

79.904

3

Li

Lithium

6.941

48

Cd

Cadmium

112.411

71

Lu

Lutetium

174.9668

55

Cs

Cesium

132.9054519

12

Mg

Magnesium

24.3050

20

Ca

Calcium

40.078

25

Mn

Manganese

54.938045

98

Cf

Californium

[251]

109

Mt

Meitnerium

[276]

6

C

Carbon

12.0107

101

Md

Mendelevium

[258]

58

Ce

Cerium

140.116

80

Hg

Mercury

17

Cl

Chlorine

35.453

42

Mo

Molybdenum

95.96

24

Cr

Chromium

51.9961

60

Nd

Neodymium

144.242

27

Co

Cobalt

58.933195

10

Ne

Neon

20.1797

29

Cu

Copper

63.546

93

Np

Neptunium

200.59

[237]

96

Cm

Curium

[247]

28

Ni

Nickel

58.6934

110

Ds

Darmstadtium

[281]

41

Nb

Niobium

92.90638

105

Db

Dubnium

66

Dy

Dysprosium

[268]

7

N

Nitrogen

14.0067

162.500

102

No

Nobelium

[259]

99

Es

Einsteinium

[252]

76

Os

Osmium

190.23

68

Er

Erbium

167.259

8

O

Oxygen

15.9994

63

Eu

Europium

151.964

46

Pd

Palladium

106.42

100

Fm

Fermium

[257]

15

P

Phosphorus

9

F

Fluorine

18.9984032

78

Pt

Platinum

87

Fr

Francium

[223]

94

Pu

Plutonium

[244]

64

Gd

Gadolinium

157.25

84

Po

Polonium

[209]

31

Ga

Gallium

69.723

19

K

Potassium

32

Ge

Germanium

72.64

59

Pr

Praseodymium

79

Au

Gold

196.966569

61

Pm

Promethium

30.973762 195.084

39.0983 140.90765 [145]

507

Atomic Number

Symbol

Name

Atomic Weight

Atomic Number

Symbol

231.03588

65

Tb

Terbium

158.92535

[226]

81

Tl

Thallium

204.3833

Name

Atomic Weight

91

Pa

Protactinium

88

Ra

Radium

86

Rn

Radon

[222]

90

Th

Thorium

232.03806

75

Re

Rhenium

186.207

69

Tm

Thulium

168.93421

45

Rh

Rhodium

102.90550

50

Sn

Tin

118.710

111

Rg

Roentgenium

[280]

22

Ti

Titanium

47.867

37

Rb

Rubidium

85.4678

74

W

Tungsten

183.84

101.07

112

Uub

Ununbium

[285]

[267]

116

Uuh

Ununhexium

[293]

44

Ru

Ruthenium

104

Rf

Rutherfordium

62

Sm

Samarium

150.36

118

Uuo

Ununoctium

[294]

21

Sc

Scandium

44.955912

115

Uup

Ununpentium

[288]

106

Sg

Seaborgium

[271]

114

Uuq

Ununquadium

[289]

34

Se

Selenium

78.96

113

Uut

Ununtrium

[284]

14

Si

Silicon

28.0855

92

U

Uranium

238.02891

47

Ag

Silver

107.8682

23

V

Vanadium

50.9415

11

Na

Sodium

22.98976928

54

Xe

Xenon

131.293

38

Sr

Strontium

87.62

70

Yb

Ytterbium

173.054

16

S

Sulfur

32.065

39

Y

Yttrium

88.90585

73

Ta

Tantalum

180.94788

30

Zn

Zinc

65.38

43

Tc

Technetium

[98]

40

Zr

Zirconium

91.224

52

Te

Tellurium

508

Appendix A

127.60

Appendix B PHYSICAL CONSTANTS Quantity Acceleration of gravity Atomic mass unit Avogadro’s number Boltzmann’s constant Charge to mass ratio of electron

Symbol or Abbreviation g amu or u NA k e/m

Value 9.80665 m/s 1.660538782 3 10227 kg 6.02214179 3 1023 particles/mol 1.3806504 3 10223 J/K 1.758820150 3 1011 C/kg

Electronic charge

e

1.602176487 3 10219 C

Electron rest mass

me

9.10938215 3 10231 kg

Faraday constant

F

96,485.3399 C/mol e2

Molar volume (STP)

Vm

22.414 L/mol

Neutron rest mass

mn

1.674927211 3 10227 kg

Planck’s constant

h

6.62606896 3 10234 J s

Proton rest mass

mp

1.672621637 3 10227 kg

Speed of light

c

2.99792458 3 108 m s21

Universal gas constant

R

8.314472 J mol21 K21 0.082057 L atm mol21 K21 62.366 L torr mol21 K21

509

Appendix C ELECTRON CONFIGURATIONS OF ATOMS IN THE GROUND STATE Z

Element

Configuration

Z

Element

Configuration

Z

Element

Configuration

1

H

1s1

37

Rb

[Kr]5s1

74

W

[Xe]4f 145d46s2

2

He

1s2

38

Sr

[Kr]5s2

75

Re

[Xe]4f 145d56s2

3

Li

[He]2s1

39

Y

[Kr]4d15s2

76

Os

[Xe]4f 145d66s2

4

Be

[He]2s2

40

Zr

[Kr]4d25s2

77

Ir

[Xe]4f 145d76s2

5

B

[He]2s22p1

41

Nb

[Kr]4d45s1

2

78

Pt

[Xe]4f 145d96s1

C

[He]2s 2p

42

Mo

[Kr]4d 5s

79

Au

[Xe]4f 145d106s1

7

N

[He]2s22p3

43

Tc

[Kr]4d 65s2

80

Hg

[Xe]4f 145d106s2

8

O

[He]2s22p4

44

Ru

[Kr]4d75s1

81

TI

[Xe]4f 145d106s26p1

5

5

1

6

2

2

8

1

9

F

[He]2s 2p

45

Rh

[Kr]4d 5s

82

Pb

[Xe]4f 145d106s26p2

10

Ne

[He]2s22p6

46

Pd

[Kr]4d10

83

Bi

[Xe]4f 145d106s26p3

11

Na

[Ne]3s1

47

Ag

[Kr]4d105s1

84

Po

[Xe]4f 145d106s26p4

12

Mg

[Ne]3s

2

13

Al

14 15

48

Cd

[Kr]4d 5s

85

At

[Xe]4f 145d106s26p5

[Ne]3s23p1

49

In

[Kr]4d105s25p1

86

Rn

[Xe]4f 145d106s26p6

Si

[Ne]3s23p2

50

Sn

[Kr]4d105s25p2

87

Fr

[Rn]7s1

P

[Ne]3s23p3

51

Sb

[Kr]4d105s25p3

88

Ra

[Rn]7s2

10

4

2

4

10

2

2

16

S

[Ne]3s 3p

52

Te

[Kr]4d 5s 5p

89

Ac

[Rn]6d17s2

17

Cl

[Ne]3s23p5

53

I

[Kr]4d105s25p5

90

Th

[Rn]6d27s2

18

Ar

[Ne]3s23p6

54

Xe

[Kr]4d105s25p6

91

Pa

[Rn]5f 26d17s2

1

92

U

[Rn]5f 36d17s2

93

Np

[Rn]5f 46d17s2

94

Pu

[Rn]5f 67s2

1

55

Cs

[Xe]6s

56

Ba

[Xe]6s2

57

La

[Xe]5d16s2

19

K

[Ar]4s

20

Ca

[Ar]4s2

21

Sc

[Ar]3d14s2 2

2

1

1

2

58

Ce

[Xe]4f 5d 6s

95

Am

[Rn]5f 77s2

22

Ti

[Ar]3d 4s

23

V

[Ar]3d34s2

59

Pr

[Xe]4f 36s2

96

Cm

[Rn]5f 76d17s2

24

Cr

[Ar]3d54s1

60

Nd

[Xe]4f 46s2

97

Bk

[Rn]5f 97s2

25

Mn

[Ar]3d54s2

61

Pm

[Xe]4f 56s2

98

Cf

[Rn]5f 107s2

99

Es

[Rn]5f 117s2

6

2

6

62

Sm

[Xe]4f 6s

2

26

Fe

[Ar]3d 4s

27

Co

[Ar]3d74s2

63

Eu

[Xe]4f 76s2

100

Fm

[Rn]5f l27s2

28

Ni

[Ar]3d84s2

64

Gd

[Xe]4f 75d16s2

101

Md

[Rn]5f 137s2

9

102

No

[Rn]5f 147s2

10

1

2

29

Cu

[Ar]3d 4s

65

Tb

[Xe]4f 6s

30

Zn

[Ar]3d104s2

66

Dy

[Xe]4f 106s2

103

Lr

[Rn]5f 146d17s2

31

Ga

[Ar]3d104s24p1

67

Ho

[Xe]4f 116s2

32

10

104

Rf

[Rn]5f 146d27s2

Er

[Xe]4f 6s

2

105

Db

[Rn]5f 146d37s2

69

Tm

[Xe]4f 136s2

106

Sg

[Rn]5f 146d47s2

[Ar]3d104s24p4

70

Yb

[Xe]4f 146s2

107

Bh

[Rn]5f 146d57s2

Br

[Ar]3d104s24p5

71

Lu

[Xe]4145d16s2

108

Hs

[Rn]5f 146d67s2

Kr

10

109

Mt

[Rn]5f 146d77s2

Ge

2

[Ar]3d 4s 4p

68

33

As

[Ar]3d104s24p3

34

Se

35 36

510

2

2

6

[Ar]3d 4s 4p

12

14

2

2

72

Hf

[Xe]4f 5d 6s

73

Ta

[Xe]4f 145d36s2

Appendix D SPECIFIC HEATS AND HEAT CAPACITIES OF SOME COMMON SUBSTANCES Substance

Specific Heat, c ( J g21 K21)

Molar Heat Capacity, Cp ( J mol21 K21)

Al(s)

0.900

24.3

Ca(s)

0.653

26.2

Cu(s)

0.385

24.5

Fe(s)

0.444

24.8

Hg(,)

0.138

27.7

H2O(s), ice

2.09

37.7

H2O(,), water

4.184

75.3

H2O(g), steam

2.03

36.4

CO2(g)

0.843

37.1

C6H6(,), benzene

1.74

136

C6H6(g), benzene

1.04

C2H5OH(,), ethanol

2.46

113

81.6

C2H5OH(g), ethanol

0.954

420

(C2H5)2O(,), diethyl ether

3.74

172

(C2H5)2O(g), diethyl ether

2.35

108

511

Appendix E SELECTED THERMODYNAMIC DATA AT 298.15 K

Species

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

Aluminum Al(s)

Species

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

Be(OH)2(s)

2907.1

50.2

2817.6

21205.20

77.91

21093.80

0

56.74

0

BeSO4(s)

0

28.3

0

AlCl3(s)

2704.2

110.7

2628.9

Al2O3(s)

21676

50.92

21582

AlPO4(s)

21733.8

90.76

21617.9

Bi(s)

Al2(SO4)3(s)

23440.84

239.3

23099.94

BiCl3(s)

2379.1

177.0

2315.0

Bi2O3(s)

2573.88

151.5

2493.7

Bi2S3(s)

2143.1

200.4

2140.6

0

5.86

0

Antimony Sb(s)

Bismuth

0

45.69

0

Sb4O6(s)

21417.1

246.0

21253.0

Boron

Sb2O5(s)

2971.9

125.1

2829.2

B(s)

SbCl5(,)

2440.2

301

2350.1

BCl3(,)

2427.2

206.3

2387.4

SbCl3(s)

2382.17

2259.4

B2H6(g)

35.6

232.11

86.7

BF3(g)

21137.00

254.12

21120.33

H3BO3(s)

21094.33

88.83

2968.92

NaBH4(s)

2183.34

104.68

2119.54

111.8

174.9

82.4

184.1

Argon Ar(g)

0

154.843

0

0

35.1

0

Br(g)

H3As(g)

66.44

222.78

68.93

Br2(,)

0

152.23

0

As2O5(s)

2924.87

105.4

2782.3

Br2(g)

30.91

245.4

3.14

AsCl3(,)

2305.0

216.3

2259.4

Arsenic As(s)

Bromine

BrF3(g)

2255.6

292.4

2229.5

HBr(g)

236.4

198.59

253.43

Barium Ba(s)

0

62.8

BaCl2(s)

2860.1

126

0 2810.9

Cd(s)

0

51.76

0

BaCO3(s)

21216.3

112.1

21137.6

CdCl2(s)

2391.50

115.27

2343.93

BaF2(s)

21207.1

96.36

21156.8

CdO(s)

2258.2

54.8

2228.4

Ba(NO3)2(s)

2992.07

213.8

2796.59

CdS(s)

2161.9

64.9

2156.5

BaO(s)

2553.5

70.42

2525.1

BaSO4(s)

21465

132

21353

Beryllium

Cadmium

Calcium Ca(s)

0

41.6

0

Ca(g)

192.6

154.8

158.9

0

9.54

0

Ca21(aq)

2542.8

253.1

2553.5

BeCl2(s)

2490.4

82.68

2445.6

CaBr2(s)

682.8

130

2663.6

BeF2(s)

21026.8

53.35

2979.4

CaC2(s)

262.8

70.3

267.8

BeO(s)

2609.6

14.14

2580.3

CaCO3(s)

21207

92.9

21129

Be(s)

512

S° (J mol21 K21)

D G f° (kJ mol21)

Species

D H f° (kJ mol21)

S° (J mol21 K21)

Species

D H f° (kJ mol21)

D G f° (kJ mol21)

CaCl2(s)

2795.0

114

2750.2

CH3NH2(,)

247.3

150.21

35.6

CaF2(s)

21215

68.87

21162

CO(NH2)2(s)

2334

105

2198

CaH2(s)

2189

42

2150

CH3CN(,)

54

150

99

CaO(s)

2635.5

40

2604.2

C3H3N(,)

172.9

188

208.6

CaS(s)

2482.4

56.5

2477.4

CO(g)

2110.5

197.6

2137.2

Ca(NO3)2(s)

2938.39

193.3

2743.07

CO2(g)

2393.5

213.6

2394.4

Ca(OH)2(s)

2986.6

76.1

2896.8

CO2(aq)

2413.80

117.6

2385.98

Ca(OH)2(aq)

21002.8

76.15

2867.6

COCl2(g)

2223.0

289.2

2210.5

Ca3(PO4)2(s)

24120.8

236.0

23884.7

CS2(g)

117.4

237.7

67.15

21433

107

21320

(CH3)2SO(,)

2203

188

299

C6H12O6(s) (glucose)

21274.5

212.1

2910.56

C12H22O11(s) (sucrose)

22221.7

360.24

21544.3

0

85.23

0

Cs (aq)

2248

133

2282

CsF(aq)

2568.6

123

2558.5

0

223.0

0

Cl(g)

121.7

165.1

105.7

Cl2(g)

2226

—

—

2

Cl (aq)

2167.29

56.48

2131.26

ClO(g)

101.22

ClO2(g)

104.60

249.4

123.4

CaSO4(s) Carbon C(s,graphite)

0

5.740

0

C(s,diamond)

1.897

2.38

2.900

C(g)

716.7

158.0

671.3

CCl4(,)

2135.4

216.4

265.27

CCl4(g)

2103

309.7

260.63

2134.5

202

273.72

CHCl3(g)

2103.1

295.6

270.37

CH2Cl2(g)

2121.46

177.8

267.26

CH3Cl(g)

280.83

234.58

257.37

CF4(g)

2925

261.61

2879

CF2Cl2(g)

2477

301

2440

CH3CF3(,)

2737

280

2668

CH4 (g)

274.81

186.2

250.75

C2H2(g)

226.7

200.8

209.2

C2H4(g)

52.26

219.5

68.12

C2H6(g)

284.86

229.5

232.9

CHCl3(,)

Cesium Cs(s) 1

Chlorine Cl2(g)

Cl2O(g)

80

266

98

ClF(g)

254.48

217.89

255.94

ClF3(g)

C3H6(g)

20.41

266.9

62.75

C3H8(g)

2103.8

269.9

223.49

C6H6(,)

49.03

172.8

124.5

C8H18(,) (n-octane)

2250.3

361.2

16.32

CH3OH(,)

2238.66

126.8

2166.27

C2H5OH(,)

2277.7

161

2174.9

C2H5OH(g)

2235.1

282.6

2168.6

HCOOH(aq)

2425.6

92.0

2351

HCOOH(,)

2424.72

128.95

2361.35

HCHO(g)

2108.57

218.77

2102.53

Cobalt

CH3CHO(g)

2166.19

250.3

2128.86

Co(s)

CH3COOH(,)

2484.5

159.8

2389.9

CoO(s)

H2C2O4(s)

2828.93

115.6

2697.2

Co3O4(s)

2891

102.5

2774

H2C2O4(aq)

2825.1

45.61

2673.9

CoCl2(s)

2312.5

109.16

2269.8

135.1

201.78

124.7

CoSO4(s)

2888.3

118.0

2782.3

HCN(g)

2163.2

281.61

2123.0

HCl(g)

292.31

186.8

295.30

HCl(aq)

2167.4

55.10

2131.2

Chromium Cr(s)

0

23.8

0

Cr2O3(s)

21139.7

81.2

21058.1

Cr2Cl3(s)

2556.5

123.0

2486.1

(NH4) 2Cr2O7(s)

21807

—

—

0

30.04

0

2237.94

52.97

2214.20

Appendix E

513

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

0

33.15

0

CuBr(s)

2104.6

96.11

2100.8

CuCl(s)

2137.2

86.2

2119.86

CuCl2(s)

2220.1

108.07

2175.7

FeCO3(s)

CuI(s)

267.8

96.7

269.5

Fe(CO)5(,)

CuO(s)

2157

42.63

2130

Cu2O(s)

2168.6

93.14

2146.0

CuS(s)

253.1

66.5

253.6

Cu2S(s)

279.5

120.9

286.2

2771.36

109

2661.8

F2(g)

2322

—

—

F2(aq)

2332.6

—

2278.8

F(g)

78.99

158.6

61.92

F2(g)

0

202.7

0

HF(g)

2271

173.7

2273

Lead

HF(aq)

2320.08

88.7

2296.8

Pb(s)

Species Copper Cu(s)

CuSO4(s)

Germanium Ge(s)

0

3

0

GeH4(g)

91

217

113

GeCl4(g)

2496

348

2457

GeO2(s)

2551

55

2497

0

48

0

Gold

Helium He(g)

0

126.150

0

218.0

H2(g)

D G f° (kJ mol21)

0

27.3

0

FeCl2(s)

2340.67

117.9

2302.3

FeCl3(s)

2399.49

142.3

2334.00

Species

114.6

Fe(s)

2741

93

2667

2774.0

338

2705.4

2733.8

445.2

2697.3

2272

—

—

Fe2O3 (s, hematite)

2824.2

87.40

2742.2

Fe3O4 (s, magnetite)

21118

146

21015

Fe(OH)2(s)

2569.0

88

2486.5

Fe(CO)5(g) FeO(s)

Fe(OH)3(s)

2823.0

106.7

2696.5

FeS2(s)

2177.5

122.2

2166.7

FeSO4(s)

2928.4

107.5

2820.8

0

64.81

0

PbCl2(s)

2359.4

136

2314.1

PbO(s, yellow)

2217.3

68.70

2187.9

PbO2(s)

2277.4

68.6

2217.33

Pb(OH)2(s)

2515.9

88

2420.9

PbS(s)

2100.4

91.2

298.7

PbSO4(s)

2919.94

148.57

2813.14

Lithium Li(s)

0

28.0

0

LiAlH4(s)

2116.3

78.74

244.7

LiCl(s)

2408.61

59.33

2384.37

LiF(s)

2615.97

35.65

2587.71

LiH(s)

290.54

20.008

268.35

203.3

LiNO3(s)

2483.13

90.0

2381.1

LiOH(s)

2487.23

50

2443.9

LiOH(aq)

2508.4

4

2451.1

Hydrogen H(g)

S° (J mol21 K21)

Iron

Fluorine

Au(s)

D H f° (kJ mol21)

0

130.6

0

H2O(,)

2285.8

69.91

2237.2

H2O(g)

2241.8

188.7

2228.6

H2O2(,)

2187.8

109.6

2120.4

Magnesium Mg(s)

Iodine

0

32.5

0

Mg21(aq)

2454.668

2138.1

2455.57

I(g)

106.6

180.66

70.16

MgBr2(s)

2524.3

117.2

2503.8

I2(s)

0

116.1

0

MgCl2(s)

2641.8

89.5

2592.3

I2(g)

62.44

260.6

19.36

MgCO3(s)

21095.8

65.7

21012.1

ICl(g)

17.78

247.4

25.52

MgF2(s)

21123.4

57.24

21070.2

HI(g)

26.5

206.5

1.72

MgI2(s)

2364.0

129.7

2358.2

514

Appendix E

Species

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

Mg(NO3)2(s)

2790.65

164.0

MgO(s)

2601.8

27

S° (J mol21 K21)

Species

D H f° (kJ mol21)

D G f° (kJ mol21)

2589.4

NH4Cl(aq)

2300.2

—

—

2569.6

NH4HCO3(s)

2847

12.1

2666

Mg(OH)2(s)

2924.7

63.14

2833.7

NH4I(s)

2201.4

117

2113

Mg3(PO4)2(s)

23780.7

189.20

23538.7

NH4NO3(s)

2365.6

151.1

2184.0

MgS(s) MgSO4(s)

2347

—

—

21284.9

91.6

21170.6

Manganese Mn(s) MnCl2(s)

(NH4)2SO4(s)

21180.85

220.1

2901.67

NF3(g)

2125

260.6

283.3

NO(g)

90.25

210.7

86.57

NO2(g)

33.2

240.0

51.30

0

32.01

0

2481.29

118.24

2440.50

N2O(g)

82.05

219.7

104.2

83.72

321.28

139.46

MnO(s)

2385.22

59.71

2362.90

N2O3(g)

MnO2(s)

2520.03

53.05

2465.14

N2O4(g)

9.16

304.2

97.82

2881.1

N2O5(g)

11

356

115

243.1

178

114

Mn2O3(s)

2959.0

110.5

Mn3O4(s)

21387.8

155.6

21283.2

N2O5(s)

MnSO4(s)

21065.25

112.1

2957.36

NOCl(g)

52.59

264

66.36

HNO3(,)

2174.1

155.6

280.79

HNO3(g)

2135.1

266.2

274.77

HNO3 (aq)

2206.6

146

2110.5

Mercury Hg(,) Hg(g)

0

76.02

0

61.32

174.96

31.82

HgCl2(s)

2224.3

146

2178.6

Hg2Cl2(s)

2224

146

2179

O2(g)

0

205.0

0

249.2

161.0

231.8

Oxygen

HgO(s, red)

290.83

70.29

258.56

O(g)

HgS(s, red)

28.2

82.4

250.6

O3(g)

143

238.8

163

OF2(g)

23

246.6

41

314.6

163.1

278.3

0

177

0

Neon Ne(g)

0

146.328

0

Phosphorus P(g) P4(s, white)

Nickel Ni(s) NiCl2(s) Ni(CO)4(g) NiO(s) NiS(g) NiSO4(s)

0

30.1

0

282.0

52.97

279.5

2602.9

410.4

2587.3

2244

38.6

2216

282.0

52.97

279.5

2872.91

92

2759.7

Nitrogen

P4(s, red)

273.6

91.2

248.5

P4(g)

58.91

279.98

24.44

PCl3(g)

2306.4

311.7

2286.3

PCl5(g)

2398.9

353

2324.6

PF3(g)

2918.8

273.24

2897.5

PH3(g)

5.4

210.1

13

P4O10(s)

22984

228.9

22698

H3PO4(s)

21281

110.5

21119

N2(g)

0

191.5

0

N(g)

472.704

153.19

455.579

NH3(g)

246.11

192.3

216.5

K(s)

0

63.6

0

NH3(aq)

280.29

111.3

226.50

KBr(s)

2393.798

95.90

2380.66

NH41(aq)

2132.51

113.4

279.31

KCl(s)

2436.5

82.6

2408.8

N2H4(,)

50.63

121.2

149.2

KClO3(s)

2391.2

143.1

2289.9

(NH4)3AsO4(aq)

21268

—

—

KClO4(s)

2432.747

82.59

2409.14

2270.83

113

2175.2

K2CO3(s)

21151.02

155.52

21063.5

2314.4

94.6

2201.5

K2Cr2O4(s)

21403.7

200.12

21295.7

NH4Br(s) NH4Cl(s)

Potassium

Appendix E

515

Species

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

Species

K2Cr2O7(s)

22061.5

291.2

21881.8

Sodium

KF(s)

2567.27

66.57

2537.75

Na(s)

0

51.0

0

108.7

153.6

78.11

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

KI(s)

2327.9

106.4

2323.0

Na(g)

KMnO4(s)

2837.2

171.71

2737.6

Na1(g)

601

—

—

2240.2

59.0

2261.9

KNO3(s) KO2(s)

2494.63

133.05

2394.86

Na1(aq)

284.93

116.7

2239.4

NaBr(s)

2359.9

86.82

2348.98

2411.0

72.38

2384

K2O2(s)

2494.1

102.1

2425.1

NaCl(s)

KOH(s)

2424.7

78.91

2378.9

NaCl(aq)

2407.1

115.5

2393.0 2254.85

KOH(aq)

2481.2

92.0

2439.6

NaClO4(s)

2383.30

142.3

K2SO4(s)

21437.79

175.56

21321.37

NaCN(s)

287.49

118.49

276.4

NaCH3COO(s)

2708.80

123.0

2607.18

21131

136

21048

Rubidium Rb(s) RbOH(aq)

0

76.78

0

2481.16

110.75

2441.24

Selenium Se(s) H2Se(g)

0

42.442

0

29.7

219.02

15.9

Na2CO3(s) NaF(s)

2573.647

51.46

2543.494

NaH(s)

256.275

40.016

233.46

NaHCO3(s)

2950.81

101.7

2851.0

NaHSO4(s)

21125.5

113.0

2992.8

NaH2PO4(s)

21538

128

21387

Na2HPO4(s)

21749

151

21609

Silicon

NaI(s)

2287.78

98.53

2286.06

Si(s)

NaNO3(s)

2467.85

116.52

2367.00

0

18.8

0

2457.3

277.8

2443.9

Na2O(s)

2414.22

75.06

2375.46

SiC(s)

265.3

16.6

262.8

Na2O2(s)

2510.87

95.0

2447.7

SiCl4(g)

2657.0

330.6

2617.0

NaOH(s)

2426.7

64.45

2379.49

NaOH(aq)

2469.6

49.8

2419.2

Na3PO4(s)

21917.40

173.80

21788.80

2364.8

83.7

2349.8

SiBr4(,)

SiCl4(,)

2687.0

239.7

2619.84

SiH4(g)

56.9

34.3

204.5

SiF4(g)

21615

282.4

21573

SiI4(g)

2132

—

—

Na2SO4(s)

21387.08

149.58

21270.16

Na2S2O3(s)

21123.0

155

21028.0

0

52.3

0

Na2S(s)

SiO2(s)

2910.9

41.84

2856.7

H2SiO3(s)

21189

134

21092

H4SiO4(s)

21481.1

192

21332.9

Na2SiO3(s)

21079

—

—

Sr(s)

H2SiF6(aq)

22331

—

—

SrCl2(s)

2828.9

114.85

2781.1

SrCO3(s)

21220.1

97.1

21140.1

SrO(s)

2592.0

54.4

2561.9

Silver Ag(s)

Strontium

0

42.55

Ag1(aq)

105.79

73.86

77.12

AgBr(s)

2100.37

107.1

296.90

S(s, rhombic)

0

31.8

0

278.8

167.8

238.3

102.30

430.98

49.63

218

331

231.8

21209

291.7

21105

AgCl(s) AgI(s)

0 Sulfur

2126.904

96.085

2109.8

S(g)

261.84

115.5

266.19

S8(g)

AgNO3(s)

2124.39

140.92

233.41

S2Cl2(g)

Ag2O(s)

231.05

121.3

211.20

SF6(g)

AgS(s)

232.59

144.01

240.67

H2S(g)

220.6

205.7

233.6

AgSCN(s)

87.9

131.0

101.36

H2SO4(,)

2814.0

156.9

2690.1

Ag2SO4(s)

2715.88

200.4

2618.41

H2SO4(aq)

2907.5

17

2742.0

516

Appendix E

Species

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

Species

SO2(g)

2296.8

248.1

2300.2

Uranium

SO3(g)

2395.6

256.6

2371.1

U(s)

SOCl2(,)

2206

—

—

UF6(s)

SO2Cl2(,)

2389

—

—

UO2(s) UO3(s)

D H f° (kJ mol21)

S° (J mol21 K21)

D G f° (kJ mol21)

0

50.21

0

22147.4

377.9

22063.7

21084.9

77.03

21031.7

21223.8

96.11

21145.9

Tellurium Te(s) TeO2(s)

0

49.71

0

Xenon

2322.6

79.5

2270.3

Xe(g)

Tin

0

169.683

0

XeF2(g)

2130

260

296

Sn(s, white)

0

51.55

0

XeF4(g)

2215

316

2138

Sn(s, gray)

22.09

44.1

0.13

XeO3(g)

502

287

561

SnCl2(s)

2350

—

—

Zinc

SnCl4(,)

2511.3

258.6

2440.2

Zn(s)

0

41.63

0

SnCl4(g)

2471.5

366

2432.2

ZnCl2(s)

2415.05

111.46

2369.398

SnO(s)

2285.8

56.5

2256.9

ZnCO3(s)

2812.78

82.4

2731.52

SnO2(s)

2580.7

52.3

2519.7

ZnO(s)

2348.3

43.64

2318.3

Zn(OH)2(s)

2643.25

81.6

2555.07

ZnS(s)

2205.6

57.7

2201.3

ZnSO4(s)

2982.8

110.5

2871.5

Titanium Ti(s)

0

30.6

0

TiCl4(,)

2804.2

252.3

2737.2

TiCl4(g)

2763.2

354.8

2726.8

0

32.6

0

2842.9

75.90

2764.1

Tungsten W(s) WO3(s)

Appendix E

517

Appendix F IONIZATION CONSTANTS OF WEAK ACIDS AT 25°C Acid Acetic Arsenic

Formula and Ionization Equation CH3COOH 1 H2O N CH3COO2 1 H3O1 2

H3AsO4 1 H2O N H2AsO4 1 H3O

Ka1 5 2.5 3 1024

H2AsO42 1 H2O N HAsO422 1 H3O1

Ka2 5 5.6 3 1028

22

1 H2O N AsO4

32

H2AsO3 1 H2O N HAsO3 Boric

B(OH)3 1 H2O N 22

BO2(OH)

1 H2O N

Butanoic

CH3CH2CH2COOH 1 H2O N

Carbonic

H2CO3 1 H2O N HCO32 1 H2O N C3H5O(COOH)3 1 H2O N C4H5O3(COOH)22 1 H2O N 22

C5H5O5COOH

1 H2O N

Cyanic

HOCN 1 H2O N

Formic

HCOOH 1 H2O N

Gluconic

HOCH2(CHOH)4COOH 1 H2O N

Heptonic

HOCH2(CHOH)5COOH 1 H2O N

Hydrazoic

HN3 1 H2O N

Hydrocyanic

HCN 1 H2O N

Hydrofluoric

HF 1 H2O N

Hydrogen peroxide Hydrosulfuric

H2O2 1 H2O N H2S 1 H2O N HS2 1 H2O N

Hypobromous

HOBr 1 H2O N

Hypochlorous

HOCl 1 H2O N

Hypoiodous Nitrous Oxalic

HOI 1 H2O N HNO2 1 H2O N (COOH)2 1 H2O N COOCOOH2 1 H2O N

Phenol Phosphoric

HC6H5O 1 H2O N H3PO4 1 H2O N 2

H2PO4 1 H2O N HPO422 1 H2O N Phosphorous

H3PO3 1 H2O N H2PO32 1 H2O N

518

Ka3 5 3.0 3 10213

1 H3O

22

1

1 H3O

C6H5COOH 1 H2O N C6H5COO2 1 H3O1 BO(OH)22 1 H2O N

Citric

1

H3AsO3 1 H2O N H2AsO32 1 H3O1 2

Benzoic

1.8 3 1025

1

HAsO4 Arsenous

Ka

BO(OH)22 1 H3O1 BO2(OH)22 1 H3O1 BO332 1 H3O1 CH3CH2CH2COO2 1 H3O1 HCO32 1 H3O1 CO322 1 H3O1 C4H5O3(COOH)22 1 H3O1 C5H5O5COOH22 1 H3O1 C6H5O732 1 H3O1 OCN2 1 H3O1 HCOO2 1 H3O1 HOCH2(CHOH)4COO– 1 H3O1 HOCH2(CHOH)5COO– 1 H3O1 N32 1 H3O1 CN2 1 H3O1 F2 1 H3O1 HO22 1 H3O1 HS2 1 H3O1 S22 1 H3O1 OBr2 1 H3O1 OCl2 1 H3O1 OI2 1 H3O1 NO22 1 H3O1 COOCOOH2 1 H3O1 (COO)222 1 H3O1 C6H5O2 1 H3O1 H2PO42 1 H3O1 HPO422 1 H3O1 HPO432 1 H3O1 H2PO32 1 H3O1 HPO322 1 H3O1

Ka1 5 6.0 3 10210 Ka2 5 3.0 3 10214 6.3 3 1025 Ka1 5 7.3 3 10210 Ka2 5 1.8 3 10213 Ka3 5 1.6 3 10214 1.5 3 1025 Ka1 5 4.4 3 1027 Ka2 5 4.8 3 10211 Ka1 5 7.4 3 1023 Ka2 5 1.7 3 1025 Ka3 5 7.4 3 1027 3.5 3 1024 1.8 3 1024 2.4 3 10–4 1.3 3 10–5 1.9 3 1025 6.2 3 10210 6.3 3 1024 2.4 3 10212 Ka1 5 1.0 3 1027 Ka2 5 1.0 3 10219 2.5 3 1029 2.9 3 1028 2.3 3 10211 4.5 3 1024 Ka1 5 5.9 3 1022 Ka2 5 6.4 3 1025 1.3 3 10210 Ka1 5 7.5 3 1023 Ka2 5 6.2 3 1028 Ka3 5 3.6 3 10213 Ka1 5 1.6 3 1022 Ka2 5 7.0 3 1027

Acid Propanoic Salicylic Selenic Selenous

Formula and Ionization Equation CH3CH2COOH 1 H2O N CH3CH2COO2 1 H3O1

1.3 3 1025

C6H4(OH)COOH 1 H2O N C6H4(OH)COO– 1 H3O1

1.1 3 10–3

H2SeO4 1 H2O N HSeO42 1 H3O1

Ka1 5 Very large

HSeO42 1 H2O N SeO422 1 H3O1

Ka2 5 1.2 3 1022

H2SeO3 1 H2O N HSeO32 1 H3O1

Ka1 5 2.7 3 1023

2

HSeO3 1 H2O N SeO3 Sulfuric Sulfurous

22

1

1 H3O

Ka2 5 2.5 3 1027

H2SO4 1 H2O N HSO42 1 H3O1

Ka1 5 Very large

HSO42 1 H2O N SO422 1 H3O1

Ka2 5 1.2 3 1022

H2SO3 1 H2O N HSO32 1 H3O1

Ka1 5 1.2 3 1022

2

HSO3 1 H2O N Tellurous

Ka

SO322

1

1 H3O

Ka2 5 6.2 3 1028

H2TeO3 1 H2O N HTeO32 1 H3O1

Ka1 5 2 3 1023

HTeO32 1 H2O N TeO322 1 H3O1

Ka2 5 1 3 1028

Appendix F

519

Appendix G IONIZATION CONSTANTS OF WEAK BASES AT 25°C Base Ammonia

Formula and Ionization Equation

Kb

NH3 1 H2O N NH41 1 OH2

1.8 3 1025

1

2

Aniline

C6H5NH2 1 H2O N C6H5NH3 1 OH

4.2 3 10210

Dimethylamine

(CH3)2NH 1 H2O N (CH3)2NH21 1 OH2

7.4 3 1024

Ethylenediamine

(CH2)2(NH2)2 1 H2O N (CH2)2(NH2)2H1 1 OH2 1

(CH2)2(NH2)2H 1 H2O N (CH2)2(NH2)2H2 Hydrazine

N2H5 1 H2O N N2H6 Methylamine Pyridine Trimethylamine

520

21

1 OH

1 OH

Kb2 5 8.9 3 10216 6.6 3 1029

2

5.0 3 1024

CH3NH2 1 H2O N CH3NH3 1 OH

C5H5N 1 H2O N C5H5NH1 1 OH2 1

Kb2 5 2.7 3 1028 Kb1 5 8.5 3 1027

2

NH2OH 1 H2O N NH3OH1 1 OH2 1

Kb1 5 8.5 3 1025 2

N2H4 1 H2O N N2H51 1 OH2 1

Hydroxylamine

21

1.5 3 1029

(CH3)3N 1 H2O N (CH3)3NH 1 OH

2

7.4 3 1025

Appendix H SOLUBILITY PRODUCT CONSTANTS OF SOME INORGANIC COMPOUNDS AT 25°C Substance

Ksp

Substance CaHPO4

Aluminum compounds AlAsO4

216

1.6 3 10

233

Al(OH)3

1.9 3 10

AlPO4

1.3 3 10220

Antimony compounds Sb2S3

1.6 3 10293

Barium compounds

Ca(H2PO4)2 Ca3(PO4)2

Ksp

Substance

2.7 3 1027

FeCO3

3.5 3 10211

233

Fe(OH)2

7.9 3 10215

FeS

4.9 3 10218

Fe4[Fe(CN)6]3

3.0 3 10241

Fe(OH)3

6.3 3 10238

Fe2S3

1.4 3 10288

1.0 3 10 2.0 3 10

Chromium compounds CrAsO4

7.8 3 10221

Cr(OH)3

6.7 3 10231

CrPO4

2.4 3 10223

Lead compounds

1.1 3 10213

Cobalt compounds

BaCO3

8.1 3 1029

Co3(AsO4)2

7.6 3 10229 8.0 3 10213

BaCrO4

2.0 3 10

CoCO3

BaF2

1.7 3 1026

Co(OH)2

2.5 3 10216

Ba3(PO4)2

229

1.3 3 10

CoS (a)

5.9 3 10221

BaSeO4

2.8 3 10211

CoS ( b)

8.7 3 10223

BaSO3

8.0 3 1027

Co(OH)3

4.0 3 10245

BaSO4

1.1 3 10210

Co2S3

2.6 3 102124

Copper compounds

Bismuth compounds BiOCl

7.0 3 1029

BiO(OH)

1.0 3 10212

BiI3

8.1 3 10219

BiPO4

1.3 3 10223

Bi2S3

1.6 3 10272

Cadmium compounds

29

1.5 3 10213

PbCl2

1.7 3 1025

PbCrO4

1.8 3 10214

PbF2

3.7 3 1028

Pb(OH)2

2.8 3 10216

PbI2

8.7 3 1029

Pb3(PO4)2

3.0 3 10244 1.5 3 1027

PbSO4

1.8 3 1028

PbS

8.4 3 10228

220

CuCN

3.2 3 10

CuI

5.1 3 10212

Cu2S

1.6 3 10248

CuSCN

1.6 3 10211

Cu3(AsO4)2

7.6 3 10236

2.5 3 10214

Cu2[Fe(CN)6]

1.3 3 10216

Cd(CN)2

1.0 3 1028

Cu(OH)2

1.6 3 10219

Cd2[Fe(CN)6]

3.2 3 10217

CuS

8.7 3 10236

214

Calcium compounds

PbCO3

PbSeO4

CdCO3

3.6 3 10

6.3 3 1026

1.9 3 1027

2.5 3 10210

CdS

PbBr2

5.3 3 10

CuCO3

229

4.1 3 10236

CuCl

2.2 3 10232

1.2 3 10

Pb3(AsO4)2

CuBr

Cd3(AsO4)2

Cd(OH)2

Iron compounds

23

Ba3(AsO4)2

210

Ksp

Gold compounds AuBr

5.0 3 10217

AuCl

2.0 3 10213

Magnesium compounds Mg3(AsO4)2

2.1 3 10220

MgC2O4

8.6 3 1025

MgF2

6.4 3 1029

Mg(OH)2

1.5 3 10211

MgNH4PO4

2.5 3 10212

Manganese compounds Mn3(AsO4)2

2.1 3 10220

MnCO3

1.8 3 10211

Mn(OH)2

4.6 3 10214

MnS

5.1 3 10215

Mn(OH)3

1.0 3 10236

Ca3(AsO4)2

6.8 3 10219

AuI

1.6 3 10223

CaCO3

4.8 3 1029

AuBr3

4.0 3 10236

CaCrO4

7.1 3 1024

AuCl3

3.2 3 10225

Mercury compounds

CaF2

1.7 3 10

Au(OH)3

1.0 3 10253

Hg2Br2

1.3 3 10222

Ca(OH)2

7.9 3 1026

AuI3

1.0 3 10246

Hg2CO3

8.9 3 10217

210

521

Substance

Ksp

Substance

1.1 3 10218

Hg2Cl2 Hg2CrO4

29

5.0 3 10

4.5 3 10229

Hg2I2

27

Ksp

AgBr

Sr3(PO4)2

1.0 3 10231

SrSO3

4.0 3 1028

213

SrSO4

2.8 3 1027

1.1 3 10 5.3 3 10

212

Hg2SO4

6.8 3 10

Ag2CO3

8.1 3 10

Hg2S

5.8 3 10244

AgCl

1.8 3 10210

Hg(CN)2

3.0 3 10223

Ag2CrO4

9.0 3 10212

Hg(OH)2

226

2.5 3 10

AgCN

6.0 3 10217

HgI2

4.0 3 10229

Ag4[Fe(CN)6]

1.6 3 10241

HgS

253

3.0 3 10

Nickel compounds

Ksp

220

Silver compounds Ag3AsO4

Substance

216

AgI

1.5 3 10

Ag3PO4

1.3 3 10220

Ag2SO3

1.5 3 10214 25

Tin compounds Sn(OH)2

2.0 3 10226

SnI2

1.0 3 1024

SnS

1.0 3 10228

Sn(OH)4

1.0 3 10257

SnS2

1.0 3 10270

Zinc compounds

Ni3(AsO4)2

1.9 3 10226

Ag2SO4

1.2 3 10

Zn3(AsO4)2

1.1 3 10227

NiCO3

6.6 3 1029

Ag2S

1.0 3 10249

ZnCO3

1.5 3 10211

Ni(CN)2

3.0 3 10223

AgSCN

1.0 3 10212

Zn(CN)2

8.0 3 10212

Ni(OH)2

2.8 3 10216

Strontium compounds

Zn2[Fe(CN)6]

4.1 3 10216

NiS (a)

3.0 3 10221

Sr3(AsO4)2

Zn(OH)2

4.5 3 10217

NiS ( b)

226

1.0 3 10

SrCO3

9.4 3 10

Zn3(PO4)2

9.1 3 10233

NiS (g)

2.0 3 10228

SrCrO4

3.6 3 1025

ZnS

1.1 3 10221

522

Appendix H

1.3 3 10218 210

Appendix I STANDARD REDUCTION POTENTIALS IN AQUEOUS SOLUTION AT 25°C Half-Reaction, Acidic Solution

Standard Reduction Potential, E° (volts)

Li1(aq) 1 e2 : Li(s)

23.045

K1(aq) 1 e2 : K(s)

22.925

1

2

Rb (aq) 1 e : Rb(s)

22.925

Ba21(aq) 1 2 e2 : Ba(s)

22.90

21

2

Sr (aq) 1 2 e : Sr(s)

22.89

Ca21(aq) 1 2 e2 : Ca(s)

22.87

1

2

Na (aq) 1 e : Na(s)

22.714

Mg21(aq) 1 2 e2 : Mg(s)

22.37

2

2

H2(g) 1 2 e : 2 H (aq)

22.25

Al31(aq) 1 3 e2 : Al(s)

21.66

41

2

Zr (aq) 1 4 e : Zr(s)

21.53

ZnS(s) 1 2 e2 : Zn(s) 1 S22(aq) 22

2

21.44

CdS(s) 1 2 e : Cd(s) 1 S (aq)

21.21

V21(aq) 1 2 e2 : V(s)

21.18

21

2

Mn (aq) 1 2 e : Mn(s)

21.18 22

2

FeS(s) 1 2 e : Fe(s) 1 S (aq)

21.01

Cr21(aq) 1 2 e2 : Cr(s)

20.91

21

2

Zn (aq) 1 2 e : Zn(s)

20.763

Cr31(aq) 1 3 e2 : Cr(s) 1

20.74 2

HgS(s) 1 2 H (aq) 1 2 e : Hg(,) 1 H2S(g)

20.72

Ga31(aq) 1 3 e2 : Ga(s)

20.53

1

2

2 CO2(g) 1 2 H (aq) 1 2 e : (COOH)2(aq)

20.49

Fe21(aq) 1 2 e2 : Fe(s)

20.44

Cr31(aq) 1 e2 : Cr21(aq)

20.41

Cd21(aq) 1 2 e2 : Cd(s)

20.403

1

2

Se(s) 1 2 H (aq) 1 2 e : H2Se(aq)

20.40

PbSO4(s) 1 2 e2 : Pb(s) 1 SO422(aq)

20.356

Tl1(aq) 1 e2 : Tl(s)

20.34

Co21(aq) 1 2 e2 : Co(s)

20.28

21

2

Ni (aq) 1 2 e : Ni(s)

20.25

[SnF6]22(aq) 1 4 e2 : Sn(s) 1 6 F2(aq)

20.25

AgI(s) 1 e2 : Ag(s) 1 I2(aq)

20.15

Sn21(aq) 1 2 e2 : Sn(s)

20.14

21

2

Pb (aq) 1 2 e : Pb(s)

20.126

N2O(g) 1 6 H1(aq) 1 H2O(,) 1 4 e2 : 2 NH3OH1(aq)

20.05

1

2

2 H (aq) 1 2 e : H2(g)

0.000

523

Half-Reaction, Acidic Solution


AgBr(s) 1 e2 : Ag(s) 1 Br2(aq)

0.10

1

2

S(s) 1 2 H (aq) 1 2 e : H2S(aq)

0.14

Sn41(aq) 1 2 e2 : Sn21(aq)

0.15

Cu21(aq) 1 e2 : Cu1(aq)

0.153

SO422(aq) 1 4 H1(aq) 1 2 e2 : H2SO3(aq) 1 H2O(,) SO422(aq)

1

2

1 4 H (aq) 1 2 e : SO2(g) 1 2 H2O(,) 2

2

0.17 0.20

AgCl(s) 1 e : Ag(s) 1 Cl (aq)

0.222

Hg2Cl2(s) 1 2 e2 : 2 Hg(,) 1 2 Cl2(aq)

0.27

21

2

Cu (aq) 1 2 e : Cu(s)

0.337

[RhCl6]32(aq) 1 3 e2 : Rh(s) 1 6 Cl2(aq)

0.44

Cu1(aq) 1 e2 : Cu(s)

0.521

TeO2(s) 1 4 H1(aq) 1 4 e2 : Te(s) 1 2 H2O(,)

0.529

2

2

0.535

I2(s) 1 2 e : 2 I (aq) H3AsO4(aq) 1 2 H1(aq) 1 2 e2 : H3AsO3(aq) 1 H2O(,) 22

22

2

2

0.58

[PtCl6] (aq) 1 2 e : [PtCl4] (aq) 1 2 Cl (aq)

0.68

O2(g) 1 2 H1(aq) 1 2 e2 : H2O2(aq)

0.682

[PtCl4]22(aq) 1 2 e2 : Pt(s) 1 4 Cl2(aq)

0.73

SbCl62(aq) 1 2 e2 : SbCl42(aq) 1 2 Cl2(aq)

0.75

31

21

2

Fe (aq) 1 e : Fe (aq)

0.771

Hg221(aq) 1 2 e2 : 2 Hg(,)

0.789

1

2

Ag (aq) 1 e : Ag(s)

0.7994

Hg21(aq) 1 2 e2 : Hg(,)

0.855

21

2

2 Hg (aq) 1 2 e :

Hg221(aq) 2

0.920

NO32(aq) 1 3 H1(aq) 1 2 e : HNO2(aq) 1 H2O(,)

0.94

NO32(aq) 1 4 H1(aq) 1 3 e2 : NO(g) 1 2 H2O(,)

0.96

Pd21(aq) 1 2 e2 : Pd(s)

0.987

AuCl42(aq)

2

2

1.00

1 3 e : Au(s) 1 4 Cl (aq)

Br2(,) 1 2 e2 : 2 Br2(aq)

1.08

ClO42(aq) 1 2 H1(aq) 1 2 e2 : ClO32(aq) 1 H2O(,) 2

1

2

IO3 (aq) 1 6 H (aq) 1 5 e :

1 — 2 I2(aq)

1 3 H2O(,)

Pt21(aq) 1 2 e2 : Pt(s) 1

1.19 1.195 1.2

2

1.229

O2(g) 1 4 H (aq) 1 4 e : 2 H2O(,) 1

21

2

MnO2(s) 1 4 H (aq) 1 2 e : Mn (aq) 1 2 H2O(,)

1.23

N2H51(aq) 1 3 H1(aq) 1 2 e2 : 2 NH41(aq)

1.24

Cr2O722(aq) 1 14 H1(aq) 1 6 Cl2(g) 1 2 e2 : 2 Cl2(aq) BrO32(aq) 1 6 H1(aq) 1 6 e2

1.33

ClO3

31

2

e : 2 Cr (aq) 1 7 H2O(,)

1.360 2

: Br (aq) 1 3 H2O(,) 1

(aq) 1 6 H (aq) 1 5 e : —2 Cl2(aq) 1 3 H2O(,)

2

1

2

Au31(aq) 1 3 e2 : Au(s) 2

1

1.44 1.47 1.50

2

21

MnO4 (aq) 1 8 H (aq) 1 5 e : Mn (aq) 1 4 H2O(,)

1.507

NaBiO3(s) 1 6 H1(aq) 1 2 e2 : Bi31(aq) 1 Na1(aq) 1 3 H2O(,)

1.6

Ce41(aq) 1 e2 : Ce31(aq)

1.61

2 HOCl(aq) 1 2 H1(aq) 1 2 e2 : Cl2(g) 1 2 H2O(,)

1.63

1

2

Au (aq) 1 e : Au(s)

524

Appendix I

1.68


Half-Reaction, Acidic Solution PbO2(s) 1 SO422(aq) 1 4 H1(aq) 1 2 e2 : PbSO4(s) 1 2 H2O(,) 1

1.685

21

2

NiO2(s) 1 4 H (aq) 1 2 e : Ni (aq) 1 2 H2O(,)

1.7

H2O2(aq) 1 2 H1(aq) 1 2 e2 : 2 H2O(,)

1.77

Pb41(aq) 1 2 e2 : Pb21(aq)

1.8

Co31(aq) 1 e2 : Co21(aq)

1.82

2

2

2.87

F2(g) 1 2 e : 2 F (aq)

STANDARD REDUCTION POTENTIALS IN AQUEOUS SOLUTION AT 25°C Standard Reduction Potential, E° (volts)

Half-Reaction, Basic Solution SiO322(aq) 1 3 H2O(,) 1 4 e2 : Si(s) 1 6 OH2(aq)

21.70

Cr(OH)3(s) 1 3 e2 : Cr(s) 1 3 OH2(aq)

21.30

[Zn(CN)4]22(aq) 1 2 e2 : Zn(s) 1 4 CN2(aq)

21.26

Zn(OH)2(s) 1 2 e2 : Zn(s) 1 2 OH2(aq)

21.245

[Zn(OH)4]22(aq) 1 2 e2 : Zn(s) 1 4 OH2(aq)

21.22

N2(g) 1 4 H2O(,) 1 4 e2 : N2H4(aq) 1 4 OH2(aq)

21.15

SO422(aq) 1 H2O(,) 1 2 e2 : SO322(aq) 1 2 OH2(aq)

20.93

2

2

Fe(OH)2(s) 1 2 e : Fe(s) 1 2 OH (aq)

20.877

2 NO32(aq) 1 2 H2O(,) 1 2 e2 : N2O4(g) 1 4 OH2(aq)

20.85

2

2

2 H2O(,) 1 2 e : H2(g) 1 2 OH (aq)

20.828

Fe(OH)3(s) 1 e2 : Fe(OH)2(s) 1 OH2(aq)

20.56

S(s) 1 2 e2 : S22(aq)

20.48

Cu(OH)2(s) 1 2 e2 : Cu(s) 1 2 OH2(aq)

20.36

CrO422(aq) 1 4 H2O(,) 1 3 e2 : Cr(OH)3(s) 1 5 OH2(aq)

20.12

MnO2(s) 1 2 H2O(,) 1 2 e2 : Mn(OH)2(s) 1 2 OH2(aq)

20.05

NO32(aq)

2

2

2

1 H2O(,) 1 2 e : NO2 (aq) 1 2 OH (aq)

0.01

O2(g) 1 H2O(,) 1 2 e2 : OOH2(aq) 1 OH2(aq)

0.076

HgO(s) 1 H2O(,) 1 2 e2 : Hg(,) 1 2 OH2(aq)

0.0984

[Co(NH3)6]31(aq) 1 e2 : [Co(NH3)6]21(aq)

0.10

2

2

N2H4(aq) 1 2 H2O(,) 1 2 e : 2 NH3(aq) 1 2 OH (aq)

0.10

2 NO22(aq) 1 3 H2O(,) 1 4 e2 : N2O(g) 1 6 OH2(aq)

0.15

2

2

Ag2O(s) 1 H2O(,) 1 2 e : 2 Ag(s) 1 2 OH (aq)

0.34

ClO42(aq) 1 H2O(,) 1 2 e2 : ClO32(aq) 1 2 OH2(aq)

0.36

2

2

O2(g) 1 2 H2O(,) 1 4 e : 4 OH (aq)

0.40

Ag2CrO4(s) 1 2 e2 : 2 Ag(s) 1 CrO422(aq)

0.446

NiO2(s) 1 2 H2O(,) 1 2 e2 : Ni(OH)2(s) 1 2 OH2(aq)

0.49

MnO42(aq) 1 e2 : MnO422(aq)

0.564

MnO42(aq) 1 2 H2O(,) 1 3 e2 : MnO2(s) 1 4 OH2(aq)

0.588

ClO32(aq) 1 3 H2O(,) 1 6 e2 : Cl2(aq) 1 6 OH2(aq)

0.62

2

2

2 NH2OH(aq) 1 2 e : N2H4(aq) 1 2 OH (aq)

0.74

OOH2(aq) 1 H2O(,) 1 2 e2 : 3 OH2(aq)

0.88

ClO2(aq) 1 H2O(,) 1 2 e2 : Cl2(aq) 1 2 OH2(aq)

0.89

Appendix I

525

Appendix J ANSWERS TO CHECK YOUR UNDERSTANDING EXERCISES Chapter 1

3.4 Molecular: Na2SO4(aq) 1 Pb(NO3)2(aq) : PbSO4(s) 1 2 NaNO3(aq)

1.1 See Figure 1.4.

Total ionic: 2 Na 1 (aq) 1 SO 4 22 (aq) 1 Pb 21 (aq) 1 2 NO32(aq) : PbSO4(s) 1 2 Na1(aq) 1 2 NO32(aq)

1.2 (a) 3, (b) 1 1.3 (a) 16.10 m, (b) 0.005 g 1.4 $299. (We ignore significant figures here since they would not usually be taken into account in calculating a price.) The amount purchased might still have to be a multiple of 5 gallons, so it might exceed the actual amount needed. 1.5 4.3 3 109 W

3.5 (a) 98.078 g mol21, (b) 63.0128 g mol21 (c) 95.0578 g mol21 3.6 70.5 mol 3.7 7 3 106 tons 3.8 C3H5N3O9 3.9 57 mol % Au, 22 mol % Ag, and 21 mol % Cu

1.6 3.34 3 104 g, or 33.4 kg 1.7

Net ionic: SO422(aq) 1 Pb21(aq) : PbSO4(s)

Steam

3.10 59 moles

Liquid water

3.11 0.216 M Chapter 4 4.1 13.5 mol H2O 4.2 3.27 g H2 4.3 9.6 g SO2 4.4 .4.5 3 103 moles H2 4.5 Same result as in the example: S8 is the limiting reactant. 4.6 20.8 g B2H6

Chapter 2

4.7 The remaining fuel is 262.2 g of aluminum.

2.1 28.1 amu 2.2 8 carbon atoms, 12 hydrogen atoms, 4 nitrogen atoms; C2H3N 2.3

4.8 80.9% yield 4.9 63.3 mL 4.10 0.828 M NaOH Chapter 5 5.1 82.1 torr

2.4 C6H9NO

5.2 310 K

2.5 (a) carbon disulfide, (b) sulfur hexafluoride, (c) dichlorine heptoxide

5.3 0.77 ft3

2.6 (a) copper(II) sulfate, (b) silver phosphate, (c) vanadium(V) oxide

5.5 6.4 g SO2, 3.4 g SO3; PSO2 5 0.98 atm, PSO3 5 0.42 atm

Chapter 3

5.6 0.96 g NaHCO3 5.7 9.67 L SO2

3.1 C3H8 1 2 O2 : 3 CH2O 1 H2O

5.8 Pideal 5 0.817 atm, PVdW 5 0.816 atm

3.2 Compounds (a), (b), and (c) are all soluble; (d) is insoluble. 3.3 Molecular : 2 HCl(aq) 1 Ca(OH)2(aq) : CaCl2(aq) 1 2 H2O(,) Total ionic : 2 H1(aq) 1 2 Cl2(aq) 1 Ca21(aq) 1 2 OH2(aq) : Ca21(aq) 1 2 Cl2(aq) 1 2 H2O(,) 1

1

Chapter 6 6.1 6.2 3 1014 s21 6.2 1.7 3 10219 J 6.3 2.6 3 10220 J

2

6.4 97.27 nm

Net ionic : H (aq) 1 OH (aq) : H2O(,) 1

(These could also be written with H3O instead of H .)

526

5.4 V 5 29.6 L, PO2 5 0.349 atm, PN2 5 1.78 atm

6.5 4d

6.6 1s22s22p63s23p2 2

10

Chapter 10

1

10.1 266.8 J mol21 K21

6.7 [Ar] 4s 3d 4p 2

14

6.8 [Xe] 6s 4 f 5d

10

6.9 F (smallest) , Si , Cr , Sr , Cs (largest) 6.10 Rb (lowest IE) , Mg , Si , N , He (highest IE) Chapter 7

10.3 8.80 J mol21 K21

7.1 The first ionization energy will be small, the second ionization energy will be slightly larger, and the third ionization energy will be extremely large. 7.2 RbBr will have the smallest lattice energy. The lattice includes only singly charged ions, and the Rb1 and Br2 ions are larger than those in KCl. 7.3 N9Cl 7.4 H

7.5

H

H

C

C

H

H

H

O

Cl O

H 7.6 O

C

O

7.8 (a) T-shape, (b) linear 7.9 (a) Trigonal planar, (b) linear

11.2 (a) Reaction is first order in H2, first order in Br2, and second order overall. (b) Reaction is first order in N2O5 and first order overall.

11.5 870 seconds 11.6 First-order reaction, k 5 0.0023 s21 11.7 1.6 3 1025 s21

8.1 Values are given in the text: 68% for body-centered cubic and 52% for simple cubic. 8.2 The dopant should be an element with more than four valence electrons. Phosphorus is the most likely choice. The band diagram will look like that in Figure 8.12. 8.3 Ne (lowest boiling point) , CH4 , CO , NH3 (highest boiling point) H

11.9 (a) Yes, it does match the actually stoichiometry and does not involve any steps in which three or more molecules collide. (b) N2O3 and NO are both intermediates in this mechanism. (c) The first two steps are unimolecular, and the third step is bimolecular. Chapter 12

Chapter 8

H

11.1 NO2: 8.0 3 1026 mol L21 s21; O2: 2.0 3 1026 mol L21 s21

11.8 42 kJ mol21

7.7 Octahedral

8.4

Chapter 11

11.4 Units for rate constants depend on the overall reaction order.

H

N

10.4 185.9 kJ mol21

11.3 No, as long as we do the calculation correctly the results will agree within any experimental error.

H H

10.2 Gases condense at low temperatures. The process is enthalpy driven because it is exothermic, but it decreases the entropy of the system. Enthalpy dominates at low temperatures.

O

N OH

Chapter 9 9.1 1223 J 9.2 16.3°C (or 16.3 K) 9.3 24.4 J mol21 °C21 9.4 64.0°C 9.5 29.3 kJ/g 9.6 23.28 3 104 J 9.7 1.9 kJ mol21 9.8 2252.6 kJ 9.9 21215 kJ (for reaction of 2 moles of CH4) 9.10 41.2 g NO, 19.2 g N2

[HNO2]2 12.1 K 5 ————————— [NO][NO2][H2O] 12.2 3 Cu21(aq) 1 2 PO432(aq) : Cu3(PO4)2(s); 1 K 5 ———————— [Cu21]3[PO432]2 12.3 Because the value of K for Mn21 is larger than that for either Ca21 or Mg21, Mn21 has a higher ability to precipitate OH– than either of the other cations. [NO]2[H2O]3 ; 12.4 (a) K 5 ——————— [NH3]2[O2]5/2

[NH3]2[O2]5/2 (b) K 5 ——————— ; [NO]2[H2O]3

[NO]4[H2O]6 (c) K 5 ——————— [NH3]4[O2]5 [NO]2 12.5 K 5 ————— and the reaction is N2(g) 1 O2(g) N 2 NO(g). [N2][O2] 12.6 [H2] 5 [I2] 5 0.002 M, [HI] 5 0.016 M 12.7 [Cl2] 5 [PCl3] 5 0.036 M, [PCl5] 5 0.004 M 12.8 Possibilities include removing HCN from the system or adding an acid (other than HCN), which would increase [H1]. 12.9 (a) Shift toward reactants, (b) no change

Appendix J

527

12.10 (a) K sp 5 [Mn 21 ][OH 2 ] 2 , (b) K sp 5 [Cu 21 ][S 22 ], (c) Ksp 5 [Cu1][I2], (d) Ksp 5 [Al31]2[SO422]3 12.11 7.7 3 1029 M; 1.0 3 1027 g/100 g H2O 12.12 3.5 3 1024 M in H2O; 7.6 3 1029 M in 0.15 M NaF 12.13 Hydrogen carbonate ion (HCO32) is the conjugate acid of carbonate ion (CO322), and hydroxide ion (OH2) is the conjugate base of water. 12.14 10.5% ionized. In this case, the assumption is introducing an error of roughly 0.6% so that whether or not it is acceptable would depend on the required accuracy. 12.15 ln K 5 236, which means that K for this reaction would be on the order of 10102. This extremely large number indicates that the complete oxidation is the very strongly preferred pathway.

13.1 Using zinc for the other electrode will give a cell potential of 1.100 V. Copper is the cathode and zinc is the anode. 22

13.2 [Ni ] 5 5 3 10

528

Appendix J

13.4 The 5-A current must flow for 1.1 3 104 s, which is about 3 hours. 13.5 0.27 V 13.6 No. The stated conditions would deposit only 100 mg of copper, so the coating will reach only about one-fourth of the desired thickness. 13.7 0.042A Chapter 14 14.1 14.2 14.3

Chapter 13

21

13.3 Reversing the reaction would change the sign of DG, giving 128.4 kJ. The positive value indicates that the reverse process is not spontaneous.

M

209 84 Po 218

4 : 205 82 Pb 1 2 He

Po

11 6C

: 115 B 1 01 b 1 n

14.4 0.283 s21; this is roughly 11 orders of magnitude larger than the decay constant for 14C. 14.5 1.1 3 104 years 14.6 7.034435 3 1010 kJ

Appendix K ANSWERS TO ODD-NUMBERED END-OF-CHAPTER EXERCISES Chapter 1 1.1 Approximately 200 g heavier. (This web exercise will have varying answers depending on the websites found.) 1.3 The first step in the scientific method is observation. Everything begins with observing something in nature and then trying to understand it. 1.5 This web exercise will have varying answers depending on the websites found. 1.7 Observations in the laboratory are normally made in the macroscopic regime of chemistry. 1.9 At the macroscopic level, distinguishing phases involves (1) determining their ability to maintain their shapes—only solids do and (2) determining if the system fills the entire volume available to it, or only part of the volume—this observation helps distinguish liquids from gases. 1.11 A picture such as this would be appropriate:

1.23 In deductive reasoning, you consider all the information (the facts) and put them together. In inductive reasoning, you infer what seems accurate and then devise ways to determine if subsequent observations fit the inferred conclusion. 1.25 Inductive reasoning. Scientists make predictions, inferences, and then seek to prove or refute the predictions by making further observations. They don’t start with a number of facts and then draw conclusions by putting them together. 1.27 Probably not. All observations that could be explained by the model conform to the predictions it makes. When this occurs, a model may inspire sufficient confidence to be labeled a theory. 1.29 Somebody asks to meet with you at 9:00 and you go in the evening when they expected you in the morning (but didn’t specify a.m.) (Other answers possible.) 1.31 Meters and amperes are base units. Grams are not (because kg is the base unit for mass). Joules and liters are both derived units. 1.33 Smallest to largest: nano-, centi-, kilo-, and giga1.35 The erg was an energy unit associated with work being done, and the calorie was associated with heat—at the time, scientists thought these two forms of energy were different things. 1.37 One million, 1 3 106

Liquid

Gas

1.13 (a) Chemical change; (b) physical change; (c) chemical change; (d) chemical change; and (e) physical change. 1.15 When steel rusts (a chemical change), it becomes more brittle and changes color. (Other answers possible.)

1.39 The upper temperature, 100, was set by body temperature. The lower temperature was set by the lowest temperature at which water would freeze (with added salt) and it was set to 0. This shows some limitations of measurement because we know that body temperature can vary over time or from one person to another. 1.41 (a) 263 K (b) 273.15 K (c) 553 K (d) 1.7 3 103 K

1.17 Looking at the picture in Exercise 1.11, the depiction of the gas shows a significant amount of empty space between particles. The empty space contributes nothing to the mass of the system even though the entire volume is occupied, so the density of the gas is lower than that of the liquid (which occupies less than the complete volume and has less space between particles).

1.43 (a) 6.213 3 101 (b) 4.14 3 1024 (c) 5.1 3 1026 (d) 8.71 3 108 (e) 9.1 3 103

1.19 If you wanted to know how many cities were in a particular state, you would need to define what constitutes a city. Is population the basis, or is the specific nature of incorporation, for example? (Other answers possible.)

1.53 125 g

1.21 Precise. The measurements might well be close to each other even though they are not close to the actual answer.

1.45 (a) 4

(b) 3

(c) 4

1.47 (a) 80.0 (b) 0.7615

(d) 4

(e) 3

(c) 14.712 (d) 0.029

1.49 0.58 g/mL 1.51 5.663 g 1.55 1.73 m, 5 feet 8.1 inches 1.57 (a) 3.47 mg (b) 0.273 mL (c) 7.25 3 1027 s (d) 1.3 3 1023 km 1.59 (a) 0.0223 L (b) 13.6 in3 (c) 0.0236 qt 1.61 523 g/L

529

1.63 4.1 L 14

1.65 8.97 3 10 L 1.67 2.35 cm 1.69 Each bullet should have a mass of 9700 mg. (Calculation gives a range of 9696–9744 mg, but the data given have just two significant figures, and therefore both ends of the range round to 9700 mg.) 1.71 In a crystal, atoms or molecules are close together and arranged in an orderly lattice. In a liquid, atoms or molecules are still fairly close together, but the arrangement is much more disordered. A picture similar to the first two panels of Figure 1.3 would illustrate the difference. 1.73 4900 cans 1.75 Aluminum has an elastic modulus of 10.0 3 106 which is less than the given number. Therefore, a frame made with the new material will seem stiffer than one made with 1aluminum.

2.3 Responses will vary. Carpet fibers, sofa fabric, drinking cups, and calculator cases are examples. 2.5 Different atoms convey different physical properties. By making polymers that contain different types of atoms, one can design polymers that have certain desirable traits. 2.7 Annual production is roughly 45 billion pounds. Depending on how usage is grouped, the top three uses are most likely (1) product containers such as shampoo bottles, (2) thin polymer films, and (3) grocery bags and trash bags. 2.9 Because the number of protons distinguishes one type of atom from another. It dictates how many electrons are present in the neutral atom and this is what makes atoms of different elements unique. 2.11 Isotopes are different atomic forms of an element that differ only in the number of neutrons in the nucleus. 2.13 (a) magnesium-24: 12 protons, 12 neutrons, 12 electrons (b) tin-119: 50 protons, 69 neutrons, 50 electrons

1.77 Compared to steel and titanium, aluminum has lower yield strength; so if strength is important to a design, aluminum would not be a good choice. 1.79 The primary difference appears to be the size of the tubes used. Aluminum frames use larger tubing than steel-frame bikes.

(c) thorium-232: 90 protons, 142 neutrons, 90 electrons (d) carbon-13: 6 protons, 7 neutrons, 6 electrons (e) copper-63: 29 protons, 34 neutrons, 29 electrons (f ) bismuth-205: 83 protons, 122 neutrons, 83 electrons 2.15 200.59 amu

1.81 This web exercise will have varying answers depending on the websites found. In addition to the material costs, the cost of welding, etc., vary from one material to another and contribute to the overall price differences.

2.17 It is the average of several values that also takes into account the frequency of occurrence of each value.

1.83 If we knew the mass of the acid solution in the battery, we could multiply that by 0.380 to find the mass of sulfuric acid present in the solution. Since we are given the density of the battery acid, looking up or measuring either the mass or the volume of the battery acid would enable us to complete the calculation.

2.21 78.72%

1.85 We could estimate the percentage of ethanol if we knew the densities of both water and ethanol. The information given allows us to find the density of the solution, and then we can set that equal to a weighted average of the densities of the two components. If we let x equal the fraction of ethanol, then we would have: rsolution 5 x(rethanol) 1 (1 – x)(rwater) Note that such a calculation assumes that the volume of the solution is the sum of the volumes of the individual liquids, but this is not necessarily true. 1.87 An object will float in water if its mass is less than the mass of water that would be displaced if the entire object were submerged. We are given dimensions for the cork and the lead cube but not masses. So we might start by looking up densities for cork and for lead. (Since the lead will be much heavier, we might neglect the mass of the cork.) If the combined mass of the cork and the lead is less than the mass of water that would occupy the combined volume of the cork and the lead, then they should float. Chapter 2 2.1 A polymer is a large molecule made of small repeating units of atoms called monomers that are bonded together. A monomer is a small group of atoms bonded together.

530

Appendix K

2.19 The percent abundance of 63Cu is 69.0%, and the percent abundance of 65Cu is 31.0%. 2.23 Cations have positive charges, whereas anions have negative charges. 2.25 (a) Na1: 11 protons, 10 electrons (b) Al31: 13 protons, 10 electrons (c) S22:16 protons, 18 electrons (d) Br2: 35 protons, 36 electrons 2.27

126 Te22 52 21

2.29 (a) Ba (b) Ti41 (c) PO432 (d) HCO32 (e) S22 (f ) ClO42 (g) Co21 (h) SO422 2.31 One common chemical used for ionic polymerization is butyl-lithium. This molecule decomposes into a lithium ion (positive charge) and a butyl ion (negative charge). An advantage of ionic polymerization is that the resulting polymer is easier to process further to induce desirable properties. 2.33 In the formula Ba(OH)2, there are one Ba atom, two O atoms, and two H atoms. 2.35 Electrons are either shared between atoms to form covalent chemical bonds or are transferred from one atom to another, creating ions that are held by ionic chemical bonds. 2.37 Molecules are groups of atoms bonded to one another with covalent bonds. Ionic compounds consist of arrays of positive and negative ions held by electrostatic attraction (ionic bonds). Therefore, the term molecule does not apply, and the formula of an ionic compound is represented by the smallest whole-number combination of ions that has no net charge, the formula unit.

2.39 Both involve lattice systems, but metals do not have alternating positive and negative ions. Instead, metals have positively charged nuclei (with some electrons) occupying the lattice points surrounded by a “sea” of mobile, negatively charged electrons. 2.41 Molecules consist of atoms that share pairs of electrons rather than having electrons donated from one atom to another (ionic compounds) or mobile across an entire lattice (metals). Such a shared pair of electrons between two atoms creates an attractive force known as a covalent bond. 2.43 Polymers exist as large numbers (hundreds, even thousands) of monomer units bonded together. It is more convenient to represent a polymer by its empirical or monomer formula. 2.45 The empirical formula for butadiene, C4H6, is C2H3. 2.47 The periodic ta ble is an organization of the elements with respect to size and structure. It allows correlating chemical and physical properties and predicting behavior trends. 2.49 It is observed that the number of hydrogen atoms with which an individual atom of various elements will combine varies cyclically from one side of the periodic table to another. This demonstrates periodicity—regular variations in the behavior of the elements. 2.51 (a) K: Group 1 (c) Ar: Group 18

(b) Mg: Group 2 (d) Br: Group 17

2.53 Sulfur and selenium are both in Group 16 so they should exhibit similar properties. 2.55 A metalloid is an element that exhibits some properties of metals and some of nonmetals. They are said to be intermediate in character and occupy the region on the periodic table between metals and nonmetals. 2.57 (a) Si: metalloid (d) N: nonmetal

(b) Zn: metal

(c) B: metalloidv

(e) K: metal

(f ) S: nonmetal

2.59 The chemistry of transition metals is more complicated because most transition metals can form multiple different cations (whose charges vary). 2.61 Organic chemistry is a field of chemistry focused on molecules that consist primarily of carbon atoms bonded to each other. 2.63 A functional group is a specific group of atoms bonded to each other that display certain properties and influences the properties of organic molecules when they are present. Organic molecules are classified and named according to which functional groups are present. 2.65 The molecular and empirical formula for the molecule on the left is C15H10N2O2. The molecular formula for the one on the right is C2H6O2; the empirical formula is CH3O. 2.67 The molecular formula for GABA is C4H9NO2. 2.69 In molecules, atoms can combine in many different ratios and the names must reflect those particular formulas. In ionic compounds, two specific ions can combine in only one way, thus the different rules for naming ionic versus molecular compounds. 2.71 (a) dinitrogen pentoxide (c) nitrogen tribromide

(b) disulfur dichloride (d) tetraphosphorus decoxide.

2.73 (a) SF6 (b) BrF5 (c) S2Cl2 (d) S4N4 2.75 (a) magnesium chloride (b) iron(II) nitrate (c) sodium sulfate (d) calcium hydroxide (e) iron(II) sulfate 2.77 (a) phosphorus pentachloride (b) sodium sulfate (c) calcium nitride

(d) iron(III) nitrate

(e) sulfur dioxide

(f ) dibromine pentoxide

2.79 Low-density polyethylene (LDPE) is made of branched molecules that cannot pack very closely together, resulting in lower density for the polymer. High-density polyethylene (HDPE) consists of straight-chain molecules that can pack together very closely, giving a higher density. LDPE has lower density, is softer and weaker; HDPE has higher density, is harder, and much stronger. 2.81 According to data from the American Plastics Council, 7.8 billion pounds of LDPE and 15.7 billion pounds of HDPE were produced in the United States in 2003. As noted earlier, HDPE is used primarily for consumer product bottles, whereas LDPE is used in industrial liners and garbage bags. 2.83 You need to use the fact that oxygen nearly always forms an oxide ion with a 22 charge in compounds with metals. Find compounds with the metal, look at their formulas, and calculate the charge on the metal that would offset the charge that is present from the oxygen atoms. In the case of iron, FeO must have Fe21 and Fe2O3 must have Fe31 for the compounds to be neutral. 2.85 The average mass is closest to the mass number of the heaviest isotope, 88Sr. This can only be true if 88Sr is the most abundant isotope. 2.87 Because oxide ions are always 22, these two compounds tell us the charge of iron ions in the compounds. FeO tells us we have a 21 ion and Fe2O3 tells us we have a 31 ion. Next we need to recall that chloride ions are always 12, so we can tell that the two common compounds must be FeCl2 and FeCl3. 2.89 Steel is mostly iron, so we will compare the densities of aluminum, iron, and titanium. The shading in Figure 2.12 shows aluminum with a lower density than the other two elements, and this is consistent with the data given. Figure 2.12 does not differentiate between the densities of iron and titanium. If more levels of shading were used, we would expect the figure to show that iron is actually denser than titanium. 2.91 The density is 0.95 g/cm3, so it is HDPE. Chapter 3 3.1 1 Release a large amount of energy. 2 Decompose into a collection of smaller, simpler molecules, usually gases. 3 Explosions occur very rapidly. 4 Generally solids. 5 A large number of gaseous molecules are formed that expand to reduce their densities—forming a shock wave. 3.3 A blast front propagates when fast moving particles near the origin of the blast collide with more distant molecules and impart some of their energy. The new molecules are now moving faster and they collide with molecules even further away and so on.

Appendix K

531

3.5 Nitrogen based explosives easily break down to generate the gases needed for an explosion, N2 is one of those gases. 3.7 s: solid; ,: liquid; g: gas; and aq: aqueous

3.37 The mole is the number of particles in exactly 12 grams of carbon-12. The mole is a convenient way for chemists to relate large numbers of molecules/atoms to one another. 1 mole 5 6.022 3 1023.

3.9 Law of conservation of matter—matter is neither created nor destroyed.

3.41 (a) Fe2O3: 159.70 g/mol

3.39 7.8 3 1019 cm3; 7.8 3 1016 L

3.11 (a) 4 Al(s) 1 3 O2(g) : 2 Al2O3(s)

(b) BCl3: 117.170 g/mol

(b) N2(g) 1 3 H2(g) : 2 NH3(g) (c) 2 C6H6(,) 1 15 O2(g) : 6 H2O(,) 1 12 CO2(g)

(c) C6H8O6: 176.126 g/mol 3.43 (a) PbS: 239.2 g/mol, 86.60% Pb, 13.40% S

3.13 2 C3H6N6O6 1 3 O2 : 6 H2O 1 6 CO2 1 6 N2

(b) C2H6: 30.069 g/mol, 19.889% C, 20.111% H

3.15 C2H5OH 1 3 O2 : 2 CO2 1 3 H2O

(c) CH3COOH: 60.052 g/mol, 40.002% C, 6.7135% H, 53.285% O

3.17 (a) N2(g) 1 3 H2(g) : 2 NH3(g)

(d) NH4NO3: 80.0432 g/mol, 34.9979% N, 5.0368% H, 59.9654% O

(b) 2 H2(g) 1 CO(g) : CH3OH(,) (c) 2 S(s) 1 3 O2(g) 12 H2O(,) : 2 H2SO4(,) 3.19 3 SiCl4 1 16 NH3 : Si3N4 1 12 NH4Cl

3.45 (a) Mg3(PO4)2: 262.8578 g/mol (b) Na2S: 78.046 g/mol

3.21 Solution: a homogenous mixture of two or more substances.

(c) N2O4: 92.0110 g/mol

Solute: the minor component(s) of a solution, the component(s) in lesser amount.

3.47 NaOH: 39.9971 g/mol, so 79.9942 g NaOH are needed

Solvent: the major component of a solution, the component found in the greatest amount.

3.51 (a) C7H5N3O6: 1.577 moles

3.49 (a) 67 g (b) 0.0698 g (c) 0.060 g (d) 1.32 3 104 g (b) CH3NO2: 1.35 moles

3.23 Concentrated solution: a solution in which the number of solute particles is high. Dilute solution: solution in which the number of solute particles is low. 3.25 An electrolyte is any solute that dissolves to produce a solution that conducts electrical current. Strong electrolytes dissociate completely, whereas weak electrolytes dissociate only partially. At similar concentrations, a solution of a strong electrolyte such as HCl will conduct electricity better than a solution of a weak electrolyte such as CH3COOH. 3.27 (a) K1, OH2 (b) K1, SO422 (c) Li1, NO32 (d) NH41, SO422 3.29 1.40 3 103 mL (b) strong base (but only slightly soluble); Mg21, OH– (c) weak acid; H1, OH–

(b) 2.8 3 1024 molecules/hr (c) 6.8 3 1025 molecules/day 3.59 C17H19NO3 3.61 (a) 40 at % Cu, 60 at % Al (b) 71 wt % NiO, 29 wt % MgO (c) 50 mol % MgO, 50 mol % FeO 3.63 (a) 5.80 M (b) 4.2 M (c) 3.41 M (d) 0.71 M 3.67 (a) 0.74 M (b) 0.0646 M (c) 4.9 3 10–3 M (d) 0.294 M 3.69 5.93 3 1022 M; 1.89 3 104 gal 3.73 C6H2N4O5 1 4 O2 : 6 CO2 1 H2O 1 2 N2

(e) strong base; Li1, OH– (f ) weak acid; H1, HSO3–, SO32– 3.33 The total ionic equation shows all species present. The net ionic equation shows only those species that undergo reaction—the spectator ions are removed. 3.35 (a) Zn(s) 1 2 HCl(aq) : H2(g) 1 ZnCl2(aq) Zn(s) 1 2 H1(aq) : H2(g) 1 Zn21(aq) (b) Mg(OH)2(s) 1 2 HCl(aq) : MgCl2(aq) 1 2 H2O(,) Mg(OH)2(s) 1 2 H1(aq) : Mg21(aq) 1 2 H2O(,) (c) 2 HNO3(aq) 1 CaCO3(s) : Ca(NO3)2(aq) 1 H2O(,) 1 CO2(g) 21

2 H (aq) 1 CaCO3(s) : Ca (aq) 1 H2O(,) 1 CO2(g) (d) 3 (NH4)2S(aq) 1 2 FeCl3(aq) : 6 NH4Cl(aq) 1 Fe2S3(s) 3 S2–(aq) 1 2 Fe31(aq) : Fe2S3(s)

Appendix K

3.57 (a) 4.7 3 1022 molecules/min

3.71 1.16 3 1021 Hg atoms

(d) strong acid; H1, Br–

532

3.55 1.70 3 1024 N atoms

3.65 (a) 0.16 mol (b) 0.0083 mol (c) 0.11 mol (d) 0.027 mol

3.31 (a) strong base; K1, OH–

1

(c) C3H6N6O6: 7.56 moles 3.53 6.8 3 1013 molecules CO2

3.75 Answers will vary depending on the sites chosen. 3.77 3 NO2 1 H2O : 2 HNO3 1 NO 3.79 4 C5H5N 1 25 O2 : 20 CO2 1 10 H2O 1 2 N2 3.81 2 NH3 1 2 CH4 1 3 O2 : 2 HCN 1 6 H2O 3.83

O 2

5 O2

2

OH 2 H2O

HO O

3.85 (a) 6.022 3 1023 octane molecules, (b) 24 moles F, (c) Z6 g C 3.87 Molecular formula is C9H12. 3.89 37% 3.91 Atom percentage and mole percentage must be the same, so the range would be 34.0–81.8 mol % Pt. 3.93 Only MgS is an electrolyte.

3.95 0.494 g KMnO4 6

3.97 1.7 3 10 gallons 3.99 4 3 1011 Cu atoms 3.101 (a) 42.2 at % Zn, 7.85 at % Cd, and 50.00 at % O (b) 42.2 mol % Zn, 7.85 mol % Cd, and 50.00 mol % O (c) 62.1 wt % Zn, 19.9 wt % Cd, and 18.0 wt % O 3.103 We know that one mole of chlorophyll contains one mole of magnesium. We can look up the molar mass of magnesium. Then we can say that the mass of one mole of magnesium must equal 2.72% of the molar mass of chlorophyll. This will let us solve the problem. 3.105 We can calculate the mass percentage of chlorine in both NaCl and MgCl2. The percentage of chlorine in the sample will be a weighted average of those two values, with each one multiplied by the fraction of its component that is actually present. As long as NaCl and MgCl2 are the only compounds present, we will be able to solve for the fraction of each. 3.107 We have the total mass of CO2 exhaled per day, and we can convert that into molecules. So all we need is an estimate of the number of exhalations per day. It should be easy to find an estimate of the average human respiration rate, which would enable us to complete the calculation. 3.109 The atomic masses of iron and oxygen are fixed, so the compound with the highest ratio of oxygen atoms to iron atoms will have the greatest mass percentage of oxygen. Looking at the formulas, it should be easy to see that this will be Fe2O3, and no further information is necessary. 3.111 Ionic bonding is expected in precipitates. 3.113 0.0762 M Fe(NO3)2 and 0.152 M NO3– 3.115 If the molar masses to be used have four or five significant figures, they will limit the accuracy of a calculation only if the other data have five or more significant figures. For a mass on the order of 100 g, this will be true if we measure masses with a resolution beyond 0.01 g. Many laboratory balances can easily do so. Chapter 4 4.1 Combustion is not always complete, and products are not always carbon dioxide and water. Moreover, gasoline is not a single compound. So describing the process with a single chemical equation is also an approximation. 4.3 In complete combustion, all carbon present is converted to carbon dioxide, whereas in incomplete combustion, some fraction of the carbon is converted only to carbon monoxide. 4.5 They are gases at room temperature, so they would evaporate from liquid gasoline. 2 mol H2 2 mol H2 1 mol O2 4.7 (a) q————— r or q—————— r or q——————r 1 mol O2 2 mol H2O 2 mol H2O 1 mol O2 2 mol H2O 2 mol H2O or q————— r or q——————r or q——————r 2 mol H2 2 mol H2 1 mol O2 2 mol H2O2 2 mol H2O2 2 mol H2O (b) q—————— r or q——————r or q——————r 2 mol H2O 1 mol O2 2 mol H2O2

1 mol O2 1 mol O2 2 mol H2O or q—————— r or q——————r or q——————r 2 mol H2O2 2 mol H2O 1 mol O2 1 mol P4 1 mol P4 1 mol P4O10 (c) q—————r or q——————r or q——————r 5 mol O2 1 mol P4O10 1 mol P4 1 mol P4O10 5 mol O2 5 mol O2 or q————— r or q——————r or q——————r 1 mol P4 1 mol P4O10 5 mol O2 2 mol KClO3 2 mol KClO3 (d) q———————r or q———————r 2 mol KCl 3 mol O2 3 mol O2 2 mol KCl or q———————r or q——————— r 2 mol KClO3 2 mol KClO3 3 mol O2 2 mol KCl r or q—————— or q—————— r 2 mol KCl 3 mol O2 4.9 (a) S8 1 8 O2 : 8 SO2 (b) 25 S8 molecules (c) 200 SO2 molecules 4.11 11.3 moles 4.13 (a) 13 moles, (b) 3.6 moles, (c) 4.9 moles, (d) 2.6 3 1023 mol. 4.15 (a) 3.6 g, (b) 3.67 g, (c) 2.5 mg, (d) 12 kg. 4.17 3.2 g Cl2 4.19 0.883 metric tons C; 5.48 metric tons Fe 4.21 910 g N2 4.23 O2 is limiting. 4.25 We must have more than 24 moles of Br2. 4.27 Conditions ARE conducive for the formation of carbon monoxide. 4.29 Sulfuric acid is the limiting reactant and forms 5.88 g SO2. 4.31 Sulfuric acid is limiting and 790 g Al2(SO4)3 forms. 4.33 Nitrogen dioxide is limiting and 3.3 kg of HNO3 forms. 4.35 SiO2 is limiting, and 11.5 g of C will remain. 4.37 Side reactions produce something other than the desired product. Sometimes the reaction itself doesn’t proceed to completion. 4.39 Reaction 1: 36.6%; Reaction 2: 1.3% Reaction 3: 34% Reaction 4: 103% 4.41 88% 4.43 150 g 4.45 66.5% 4.47 H2S is limiting and the percentage yield is 79%. 4.49 0.302 g H2 4.51 It changes color when the chemical being added has consumed the chemical of interest in the solution. 4.53 22.9 mL 4.55 3.61 g 4.57 Octane ratings measure the quality of a gasoline in terms of the extent to which it can be compressed before it ignites without a spark. Octane itself is rather compressible, so if you had pure octane the rating would be 100. It is possible to have an octane rating greater than 100 for a fuel that is more compressible than octane.

Appendix K

533

4.59 Web question that will vary depending on location. 4.61 Lead reacts with metals such as rhodium, palladium, and platinum, effectively covering these metals and eliminating their catalytic activity. 4.63 104 g/mol 13

4.65 2.0 3 10 g Ca5(PO4)3F 4.67 0.031 M 4.69 5 3 108 kg NH3 4.71 Yes, CO was a limiting reactant. CO 1 2 H2 : CH3OH 4.73 Sulfur is limiting and 67.4 g of S2Cl2 forms. 4.75 HF is limiting and 1049 kg of Na3AlF6 forms. 4.77 HCl is limiting and 1.5 g CO2 forms. The concentration of Ca21 is 0.050 M. 4.79 588.2 g 4.81 CaCO 3 and Ca(OH) 2 could each neutralize two moles of HCl for each mole of base, whereas NH3 and NaOH could only neutralize one mole of HCl for each mole of base. Since the costs are given in terms of mass, we also need to consider the molar masses. Combining the reaction stoichiometry with the molar masses allows us to calculate the cost per mole of acid to be neutralized, and the results indicate that NH3 offers the lowest cost. 4.83 We are given the mass of Cr2O3, and we can convert that into moles. We can see from the balanced equation that we need two moles of Al for each mole of Cr2O3, and this information will enable us to find the number of moles of Al needed. That can be converted to mass if desired; no additional information should be needed. 4.85 The increase in mass is due to oxygen, so we know that 0.5386 g of the unknown element combines with 0.1725 g of oxygen. We can convert that mass of oxygen into moles and then use the 3:2 mole ratio from the formula to find the number of moles of element X in the original sample. Dividing the mass by the number of moles will give the desired molar mass. 4.87 There are several ways we could approach this problem. One would be to calculate the mass percentage of selenium in Bi2Se3. That would allow us to find the mass of selenium in 5 kg of Bi2Se3. Comparing that mass to the 560 g available would then tell us whether or not we have enough selenium. 4.89 We can convert 500 g O2 into moles and then use the 2:1 mole ratio from the balanced equation to find the number of moles of KMnO4 needed. 4.91 Convert 150 tons of Freon-12 to grams and then to moles. Use the 2:1 mole ratio from the balanced equation to find the number of moles of sodium oxalate needed to react with all of the Freon-12. The number of moles of sodium oxalate needed could then be converted to mass if desired.

5.3 Because NO is being consumed by reactions that are involved in the formation of photochemical smog. The rates of the reactions that consume NO exceed those that form it during daylight hours. 5.5 No, it would not contribute significantly to the formation of smog. The organic chemicals present in asphalt are not volatile, so they don’t fulfill the same role as VOCs in the formation of smog. 5.7 Gases become less dense as their temperature increases. The air in a hot-air balloon is at a higher temperature than the surrounding air. The result is that it is less dense and therefore floats on the denser, cooler air. 5.9 Atmospheric pressure is created when the mass of the gases attracted to the surface of the earth by gravity generates a force that produces pressure (force/unit area). 5.11 For a liquid with half the density of mercury, atmospheric pressure will generate a column twice as high as that generated for mercury. The column will be 750 mm 3 2 5 1500 mm. 5.13 1.5 3 102 atm 5.15 (a) 0.9474 atm (b) 9.50 3 10 2 mm Hg (c) 542 torr (d) 98.7 kPa (e) 6.91 atm 5.17 The increased pressure due to the weight of the water compresses the air in your ear, which can cause a “pop” in your ear upon expansion. 5.19 113 mL 5.21 A balloon will return to its shape because volume is proportional to absolute temperature. At constant pressure and moles of gas, lowering T will decrease the volume of the balloon, but returning to room temperature will restore the shape. 5.23 Jacques Charles first observed that when pressure and moles of gas are held constant, volume and temperature vary inversely. This relationship is a straight line when T is plotted vs. V. Although it is not possible to lower the temperature of all gases to absolute zero, if the lines that are generated by different gases are extrapolated, all lead to the same temperature and it is absolute zero (or 2273°C). 5.25 At constant volume, the pressure of a gas is directly related to absolute temperature. Increasing temperature could cause dangerously high pressure and possible tank rupture. 5.27 6.9 mL 5.29 1.38 3 103 L 5.31 61.7 g/mol 5.33 (a) 5.39 g/L (b) 1.96 g/L (c) 1.63 g/L 5.35 670 K 5.37 Partial pressure is the pressure exerted by a gas in a mixture of gases, behaving as if it were alone in the container.

4.93 C5H3

5.39 Pi 5 Ptotal Xi

4.95 91.6% Al

5.41 0.72 atm

Chapter 5

5.43 1.9 3 1023 mol NO2 and 3.3 3 1023 mol SO2

5.1 Nitrogen oxides and VOCs (volatile organic chemicals). Nitrogen oxides are emitted from automobile exhaust as are small amounts of VOCs. Gasoline vapors and spills and chemical plants also release VOCs.

534

Appendix K

5.45 PCO2 5 640 torr, PNO 5 37 torr, PSO 5 5.7 torr, and 2 PH2O 5 70. torr 5.47 Total pressure is 20.0 atm: PH2O 5 6.28 atm, PO2 5 2.91 atm, and PN2 5 10.8 atm.

5.49 Web question—answers will vary depending on the sites found by students. 5.51 0.10 g NH4Cl 5.53 1.07 g Zn 5.55 160 L O2 5.57 1.00 atm 5.59 Ammonia is the limiting reactant and 8.0 g urea form. 5.61 0.52 L 5.63 74 g NaN3 5.65 187 torr 5.67 3 3 10217 atm 5.69 High pressure and/or low temperature 5.71 CH4, N2, Kr, CH2Cl2 5.73 At higher temperature, the distribution shifts toward higher molecular speeds, so the blue curve represents 1000 K. 5.75 Mean free path is the average distance a particle travels between collisions with other particles. 5.77 (a) 16.3 atm (b) 15.8 atm 5.79 Statements (a), (b), (c), and (d) are true. Statement (e) is false. 5.81 An electric current is the target of the measurement. Because the number of electrons is related to the number of molecules of gas, the current can be converted into pressure by counting the number of molecules (via the current they produce upon ionization).

convert between moles of carbonate and moles of carbon dioxide. Finally, the volume can be determined using the ideal gas law, as long as the temperature and pressure are known or measurable. 5.101 N2: 75.518% O2: 23.15% Ar: 1.29% CO2: 0.048% 5.103 (a) 2 3 1027 molecules (b) 4 3 1021 molecules 5.105 C2H2F4 Chapter 6 6.1 Both molten iron and the filament in a light bulb are metals. When metals are heated, they emit light. The hotter the metal, the whiter the light becomes. The emitted light starts out red and progresses to orange and then to white as the temperature increases. 6.3 Tungsten is the metal used for the filament in most light bulbs. Tungsten is chosen because it can be heated to white heat without melting. 6.5 An electric arc is an electrical discharge that passes through a material that does not conduct electrical current well. 6.7 When the term light is used, normally it refers to visible light. Like other forms of radiation, visible light is composed of an electrical field that lies perpendicularly to a magnetic field—and thus it is called electromagnetic radiation. 6.9 Radio, IR, visible, UV 6.11 480 nm: 6.2 3 1014 Hz green light

5.83 Responses will vary depending on the types of gauges found.

530 nm: 5.7 3 1014 Hz yellow light

5.85 2.19 g/L

700 nm: 4 3 1014 Hz red light

5.87 34.8 torr 5.89 79% Zn 5.91 3 3 104 L air 5.93 (a) 4 Fe21(aq) 1 4 H1(aq) 1 O2(g) : 4 Fe31(aq) 1 2 H2O(,) (b) 1.4 3 103 g H2O (c) 5.1 atm 5.95 539 hours 5.97 For this problem, you need to determine whether air is present by measuring various densities and determining if there are differences. First, the unopened can must be weighed and its volume determined (by displacement if possible). Next, the soda must have its density determined separately, perhaps by measuring the mass of a volumetric container such as a graduated cylinder when empty and when filled with a known volume of soda. There’s some chance that the density of the aluminum can be ignored in this problem, but if necessary, it could be measured. If the density measured for the whole system varies from that measured for the soda alone, the most logical source for the difference would be gas in the “head space.” 5.99 To carry out this problem requires first determining the amount of carbonate present (by multiplying the amount of ore by the fraction of carbonate). A stoichiometry problem can then be set up using the chemical equation to

580 nm: 5.2 3 1014 Hz orange light 6.13 Photon is the term used to refer to a discrete amount of energy imparted by light. It is sometimes thought of as a packet of electromagnetic radiation that carries energy. 6.15 (a) 1.01 3 10220 J (b) 1.1 3 10215 J (c) 4.12 3 10223 J 6.17 (a) 5.7 3 1026 m (IR) (b) 2.3 m (radio) (c) 2.8 3 1029 m (X-ray) (d) 36 m (radio) 6.19 (a) 2.5 3 10–19 J (b) 3.1 3 10–19 J (c) 4.90 3 10–19 J 6.21 (a) 3.00 3 10–3 m (b) 6.63 3 10–23 J/photon (c) 39.9 J/mol 6.23 7.1 3 10219 J 6.25 A continuous spectrum has no missing frequencies of light, similar to a rainbow or the light given off by the Sun. A discrete spectrum consists of specific, wellseparated frequencies, like the emission spectra of elements. 6.27 The Bohr model depicted electrons circling the nucleus in orbits that could only be at certain discrete distances from the nucleus. Because energy is related to the distance between charged particles, these orbits corresponded to energy levels. For the excited electron in the Bohr atom to return to the original orbit, the atom releases a photon of energy equal to the energy difference between the two orbits. The spectra of hydrogen atoms therefore correspond to the specific energies released when electrons drop from higher energy orbits to lower energy ones.

Appendix K

535

6.29 The ground state occurs when all electrons in an atom are in the lowest energy states possible. 6.31 (a) 4.558 3 10219 J for the 435.8 nm transition and 3.638 3 10219 J for the 546.1 nm transition. (b) 1.389 3 1014 Hz. 6.33 The mercury in fluorescent bulbs emits UV light (we cannot see UV light). The UV light is absorbed by the phosphors on the inner surface of the fluorescent bulb and then is reemitted as visible light. 6.35 The wave function is used to describe electrons around the nucleus of an atom because electrons can be diffracted. Diffraction is usually associated with electromagnetic radiation (light), but electrons also produce diffraction patterns. This means that electrons also exhibit particle-wave duality. In some instances, electrons are treated as particles; in other instances, they are treated as waves. 6.37 The principal quantum number, n, has integral values starting at 1 and continuing to infinity. The secondary quantum number, ,, is limited in value by the principal quantum number and has integral values from 0 through n 2 1. 6.39 4d 6.41 (a) 8 (b) 6 (c) 2 (d) 1 6.43 An s orbital is a sphere, which means it can be rotated by any number of degrees and its appearance does not change. There are three p orbitals in a subshell, and these orbitals must be arranged so that the electrostatic repulsion from the electrons in the orbitals is minimized. This occurs when the p orbitals are arranged along the x, y, and z axes—or in specific directions. The same type of reasoning is true for the 5d orbitals, though they don’t point along the x, y, and z axes. y

6.45

x

6.47 Spin paired means that the two paired electrons in the same orbital have opposites spins, 11/2 and 21/2, or f. 6.49 Shielding describes the effect that the inner core electrons have on the nuclear charge that the outer, valence electrons experience. The inner core electrons partially mask the nuclear charge on the nucleus from the valence electrons. 6.51 The effective nuclear charge is the apparent nuclear charge that the valence electrons experience as the result of shielding by the inner core electrons. As the effective nuclear charge increases, the electrons contained within an orbital feel a stronger attractive force toward the nucleus. The electrons are pulled toward the nucleus, reducing the size of the atomic orbital. 6.53 Hund’s rule states that for a set of equivalent atomic orbitals, each orbital is filled with one electron before pairing is allowed to occur. When two electrons are placed in the same atomic orbital, electron repulsion occurs. The atomic orbital is now at a higher energy than before (repulsion increases energy). The lowest energy state is the most stable

536

Appendix K

state for an atom, so each atomic orbital in a set of equivalent orbitals receives one electron until all orbitals contain one electron, so that the atom remains at the lowest possible energy level. 6.55 Ground state configurations: (a) and (b) Excited state configurations: (c), (e), and (f ) Impossible configuration: (d) 6.57 The electronic configurations for F, Cl, and Br are listed below. F 5 1s22s22p5 or [He] 2s22p5 Cl 5 1s22s22p63s23p5 or [Ne] 3s23p5 Br 5 1s22s22p63s23p64s23d104p5 or [Ar] 4s23d104p5 The valence electron configuration of each is identical: ns2 np5 (the d electrons in Br are not valence electrons because only the outermost s and p electrons and electrons in unfilled d orbitals are considered valence electrons). Because these three halogens have similar valence electron electronic configurations, they will have similar chemistries (the valence electrons are where chemistry occurs). For example, each halogen will accept one electron to form an anion, F2, Cl2, and, Br2. When paired with Na1, they form sodium salts, NaF, NaCl, and NaBr. 6.59 All elements in a group of the periodic table have the same basic electron configuration for their valence electrons. These valence electrons are responsible for the chemistries of the elements. Because the group elements have similar configurations, they will have similar chemical properties. For example, oxygen, sulfur, selenium, and tellurium have the same valence electronic configuration, ns2 np4. They react with hydrogen to form H2O, H2S, H2Se, and H2Te. 6.61 The 7p block would be completely filled; there would be an eighth s row, and a g block (8g) would be added to the periodic table. The f and d blocks would increase by one row, as would the eighth row in the p block. This assumes that the last current element is Uuq, Z 5 114. 6.63 A halogen lamp contains halogen vapor, which reacts with the vaporized tungsten from the filament. The tungsten halogen compound then reacts with the filament, depositing the evaporated tungsten back on the filament. The result is that the filament lasts longer and can be heated to a higher temperature, generating more light. 6.65 As the effective nuclear charge increases across the periodic table, the size of the atomic orbitals decreases. The decrease in the size of the atomic orbitals results in a smaller atomic size. 6.67 (a) Be, Li, Na (b) F, N, P (c) O, I, Sn 6.69 The first large jump in ionization energies in Cl would occur between the fifth and sixth ionization energy. The fifth ionization energy would remove the last 3p electron. The sixth ionization would remove the first 3s electron. It is harder to remove electrons from a completely filled set of orbitals than it is from a set of partially filled orbitals. The next large jump in ionization energy would occur between removing the last 3s electron and the first 2p electron (the seventh and eighth ionization energies).

6.71 The electron configuration for oxygen is 1s22s22p4. The electron configuration for nitrogen is 1s22s22p3. For oxygen, there is one set of paired electrons in the 2p orbitals. There are no paired electrons in the 2p orbitals for nitrogen. The paired electrons in oxygen repel each other. The electron repulsion raises the energy of these paired 2p electrons in oxygen, reducing the amount of energy needed to remove one of the paired electrons to form a cation, compared to nitrogen. 6.73 (a) Se (b) Br– (c) Na (d) N (e) N3– 6.75 (a) Na (b) C (c) Na, Al, B, C

be chosen such that polymerization will proceed at a reasonable rate at body temperature. 7.5 Coatings can be used to improve biocompatibility. Reactants used to generate such coatings are called precursors. 7.7 Because removing a second electron from sodium ion, Na1, involves a core electron. The energy needed to accomplish this is quite large, so it is not observed in nature. 7.9 (a) N (b) Ba21 (c) Se (d) Co31 7.11 (a) Cl–, S2–, P3– (b) O2–, S2–, Se2–

6.77 778.0 nm

(c) N3–, S2–, Br–

6.79 1.3 3 109 J

(d) Cl–, Br–, I–

6.81 B

A

C

7.13 Because the ions are closely packed in an ionic solid, the next nearest neighbors are close enough to contribute their repulsive interactions to the overall lattice energy (and so on). 7.15 In a covalent bond, the electrons involved are shared between the nuclei, whereas in ionic compounds, an electron is transferred from one atom to another. Ultimately the difference lies in the extent to which an electron is transferred during bond formation.

A is the absorption event at 185 nm, and B is the emission event at 6.88 3 1014 Hz. C is the energy difference between the ground state and the final excited state. The ground state is assumed to have an energy value of zero. EC 5 6.2 3 10219 J 6.83 The unknown metal must be lithium.

Energy

7.17

Bond distance 0 Bond energy

6.85 (a) Ga31, Ca21, K1 (b) Be21, Mg21, Ca21, Ba21 (c) Al31, K1, Sr21, Rb1 (d) Ca21, K1, Rb1 6.87 n 5 5.3 3 1013 s–1, l 5 5.7 3 10–6 m 6.89 3 3 103 W, 1 3 1022 photons/s 6.91 The key features of the experiment is that the manganese atoms absorb light near 403 nm. The amount of light absorbed will be related to the number of manganese atoms present. So, if the observed absorption can be calibrated against some standard for concentration, the comparison will tell us the amount of manganese in the sample. 6.93 We are using inductive reasoning. If we assume that any energy level is possible, we would infer that the spectrum would be continuous, which disagrees with observations. Therefore, we have used inductive reasoning to say that the energy levels must be specific. 6.95 (a) 1.91 3 1027 m (b) 6.4 3 1018 photons Chapter 7 7.1 Biocompatibility is the ability of a material to coexist with a natural biological system into which it has been implanted. 7.3 PMMA bone cement works somewhat like a grout, filling space between bone and a surgical implant. A putty containing liquid methyl methacrylate and other materials is applied by the surgeon and then hardens in place as the MMA polymerizes. The composition of the cement must

Distance

7.19 The negatively charged electrons are attracted to both positively charged nuclei, so the forces involved are still coulombic. 7.21 The number of bonds that can be formed is limited by the number of openings each element has in the valence shell. It’s possible to get slightly more bonding than dictated by this number (commonly called the valence) in some cases, but it’s not possible to keep going forever. From a physical perspective, the repulsion that would arise from packing too much electron density together forbids the formation of too many bonds. 7.23 Lt 7.25 A lone pair is a pair of electrons located only on a single atom, not shared between atoms. 7.27 A double bond is stronger than two single bonds but not twice as strong. This leads to some reactivity trends where the “second” bond is broken and replaced with a new single bond (going from a double bond to two single bonds). 7.29 Electron affinity is defined for an isolated atom and measures the energy change for adding an electron. Electronegativity is defined within chemical compounds and it estimates the propensity of an atom to attract electrons to itself.

Appendix K

537

7.31 Higher electronegativity

Lower electronegativity partial + charge

partial – charge

7.51 PH3 is the one with a single lone pair. F2 has three lone pairs on each fluorine atom, CH4 has no lone pairs, and the carbonate ion has lone pairs on each oxygen atom. 7.53 (a) There are too many valence electrons shown. O3 should have only 18 valence electrons. (b) There are too few valence electrons shown. There should be 36 and only 34 are shown

If we allow the density of electrons to be shown by shading, there is greater electron density near the nucleus on the left and that portion of the bond would have a negative charge. 7.33 Such a bond is called polar because the asymmetrical distribution of electrons results in a small electric dipole (or field). The existence of poles in this field is the origin of the terminology calling it a polar bond. 7.35 (a) O!H is the most polar; C!H and S!H have very similar polarity because they have the same difference in electronegativity. (b) H!Cl is the most polar bond and Cl!Cl is the least (nonpolar). (c) C!F is the most polar and F!F is the least (nonpolar). (d) N!H is the most polar and N!Cl is the least polar.

7.55 Wave interference is the buildup or diminishment of the amplitude of a wave because of the interaction of two or more waves. Buildup results from constructive interference, and diminishment results from destructive interference. 7.57 Overlap is what is needed for constructive or destructive interference of electron waves. If electrons didn’t behave as waves, they wouldn’t interfere. So the wave behavior of electrons implies the importance of overlap. 7.59 Both sigma and pi bonds build electron density in the space between nuclei (so that is how they are similar). They differ based on the precise location of the electron density. For a sigma bond, it builds up on the line between nuclei, whereas for a pi bond, the built-up density is above and below that line. 7.61 (a) S

(c)

(d) Cl

F F S F

F F F 7.43 (a) F

C

Cl

Cl

(b) H

(c) O

(d) O

O

N

S

O

7.63 Molecular shapes provide a motivation to consider an extension to the concept of atomic orbitals combining to form molecules. For a molecule such as water, if only p orbitals are available, the bond angle would have to be 90°. But that is not the observed angle, so we must have additional flexibility, and hybrid orbitals provide that. 7.65 This type of hybrid is called an sp3 hybrid. Because four atomic orbitals are used, there must be four hybrid orbitals as well.

F H

7.71 (a) T-shape (b) Trigonal pyramid (c) Seesaw (d) Tetrahedral

C

7.73 XeF2, CO2, and BeF2 are linear. XeF2 and OF2 have lone pairs on the central atom.

7.45 In Lewis structures, resonance occurs when there are two or more equivalent ways to draw a double bond. The actual structure is an average of all of the equivalent structures. 7.47 It is an average of all of the equivalent structures. In plants, crossbred varieties are referred to as hybrids and they are an average of the two original plants. 7.49 When resonance structures exist so that there is a hybrid, the electrons involved in that component of the molecule are not localized to the same small space commonly associated with a chemical bond. Electrons with “more space” will have lower energy than more localized electrons.

Appendix K

Cl

7.69 Like charges repel each other. Clouds of electron density will all be negatively charged, so they will tend to repel each other.

N

538

H

7.67 (a) sp3 (b) sp3 (c) sp (d) sp2

C

H

C

F C

F

C

N

H

(b)

S

H

7.37 No, they cannot. Since fluorine is the most electronegative element, no other element can pull electron density away from a fluorine atom. 7.39 The most polar diatomic would be At!F because it has the greatest difference in electronegativity. (Astatine is radioactive. For nonradioactive elements, I!F would be the most polar.) 7.41 (a) :C9O: (b) H S H

C

This carbon is trigonal planar.

7.75

H

H C

C C

H This carbon is trigonal planar.

H

H

H

This carbon is tetrahedral.

7.77 The geometry about the C"C double bond is trigonal planar, but when the reaction occurs, the geometry becomes tetrahedral. 7.79 The honeycomb structure shown in Figure 7.18 means that nearly all of the SiO4 groups are at a surface.

In an ionic solid, oppositely charged ions are held in place and cannot move to conduct electricity. In pure water, there are no charges that could move. When an ionic solid dissolves, however, the charges are now free to move and electricity can be conducted. Any molecular level picture has to show these features.

7.81 Pores are filled with small drug molecules Pores are capped with materials that can be triggered by something characteristic of a disease

7.97 Because all have two ns and two np electrons in their valence shells, they show four dots. 7.99 sp3 hybridization in product

sp2 hybridization in reactant

Empty pores do not have the cap and have dispensed their drug molecules

7.83 (a)

(b)

F

O

H

C

C

H

H

H

H C

H

R

C H

O

H

H

H

H

C

C

C

C

H

H

H

H

Chapter 8

F F

F

F

8.1 There are three solid forms of carbon: diamond, graphite, and buckminsterfullerene (and other forms of fullerenes and nanotubes.)

sp 3

sp 3d 2 Square planar

8.3 For engineering applications, the hardness of diamonds leads to important applications. If you consider a material in which you wish to drill a hole, the material you use in your drill bit must be harder than the substance. Because diamond is so hard, it can drill through most substances— even rather hard ones—so it has useful applications

F

Si

F

Tetrahedral

(c)

F

Kr

F F

Se

F

F sp 3d Seesaw

7.85 Based on the valence of each atom, a neutral molecule would have 32 electrons. There are 34 electrons shown in the structure, so it must have a 2- charge. Each carbon has one double bond, so the hybridization must be sp2. The structure has five sigma bonds and two pi bonds. 7.87 Cl

R

H

M

Cl

Cl

8.5 Carbon nanotubes have “caps” on their ends. These ends close off the cylindrical nanotube. The caps are one hemisphere of a molecule of buckminsterfullerene. 8.7 Answers will vary slightly based on websites found, but they should note that the atmosphere needs to be inert. If an arc is conducted in air, the oxygen present will react with the carbon and buckminsterfullerene will not form. 8.9 The packing efficiency tells us the amount of space occupied by particles in a solid. The more space filled, the more efficient the packing. 8.11 Lower efficiency

One metal that would make sense as M is bismuth, Bi. It is in the same group as N, which would produce a similar structure. Since bismuth is the only metal in that group, it is the best choice. 7.89 Because it is linear, a CO 2 molecule is nonpolar. Even though each C5O bond is somewhat polar due to the electronegativity difference, the dipole vectors from the two bonds cancel one another. 7.91 No, lead selenide should not be considered ionic. According to Figure 7.8, the electronegativities are 2.4 for selenium and 1.7 for lead. A difference of just 0.7 is not large enough to give rise to ionic bonding.

Higher efficiency

7.93 The Lewis structure of nitrogen gas shows that it has a triple bond, :N9N: Such a bond is quite strong, so it is relatively difficult to break. This makes nitrogen a relatively inert gas at normal temperatures. The triple bond can be broken, however, at higher temperatures such as those in internal combustion engines. 7.95 The key feature that would need to be in any picture is that for a current to flow, a charge must be able to move.

Appendix K

539

8.13 For the square, it is 78.5%. For the hexagonal, it is 90.7%. 8.15 The coordination number is 4. Each carbon atom is bonded to four other carbon atoms. 8.17 r 5 1.67 3 1028 cm 8.19 a 5 4.76 3 1028 cm 8.21 From the data given, Avogadro’s number must be 6.023 3 1023, a value fairly close to the accepted value.

8.41 The extent to which a molecule responds to an external electric field is measured by its polarizability. So a material that is not very polarizable will not respond very strongly to such a field. 8.43 For a substance to be a gas, its particles must not interact very strongly. Small molecules often have relatively small interactions. Large polymer molecules experience strong attractive forces and exist only in condensed phases.

8.23 Malleability requires that the metal remain bonded despite being deformed. The sea of electrons model accounts for this property because the electrons can move rapidly to adjust to changes in the positions of the positively charged cores. Changing the shape of the metal changes the positions of these cores, but it doesn’t affect the bonding significantly.

8.45 When a polar solvent is present, the dipole on the solvent molecules can interact with the electric charge on an ion.

8.25 To conduct electricity, the material must allow charge to move. In the sea of electrons model, the electrons in the “sea” can respond to an electric field by moving. So metals conduct electricity.

(c) Dispersion and dipole-dipole forces are both present. Dispersion may be more important due to the large size of the chlorine atoms.

8.27 In a bonding orbital, there is a buildup of electron density between the nuclei being bonded. In an antibonding orbital there is a node, a region where the electrons cannot be, between the nuclei involved.

(e) hydrogen bonding

8.29 The picture looks like this.

Energy

Conduction band

Valence band

8.47 All have high electronegativity. 8.49 (a) dispersion (b) hydrogen bonding

(d) dispersion (f ) dipole-dipole 8.51 Carbon tetrachloride is rather large for a molecule with only five atoms because the chlorine atoms are fairly large. Large molecules tend to be polarizable, so the dispersion forces in CCl4 are large enough that it is a liquid at room temperature. Ammonia, with only nitrogen and hydrogen has small dispersion forces, so its interactions are dominated by dipole–dipole and hydrogen bonding interactions. The fact that CCl4 is a liquid and NH3 is a gas implies that the large dispersion forces in the former are stronger than the hydrogen bonding and dipole forces in the latter.

8.31 Semiconductors have filled valence bands, but the conduction band is nearby in energy, so that some electrons occupy it by thermal excitation. As temperature is increased, thermal excitation increases, so a higher temperature will tend to increase the conductivity of the semiconductor.

8.53 For molecules to leave a liquid and enter the gas phase, they must overcome interactions with other molecules in the liquid. The stronger these interactions, the less likely it is that molecules will escape the liquid. So, the vapor pressure will be lower.

8.33 Doping either provides additional electrons that can be part of the conduction band, or it creates “holes” that also can be used to increase the conductivity of the material.

8.55 (a) increases

8.35 A semiconductor is named p-type because a dopant creates “positive” holes in the valence band. To do this, the dopant must have fewer valence electrons than the atom that is replaced. 8.37 For a voltage-gated device, the flow of charge across some plane (in this case a junction between the two types of semiconductors) is regulated by a change in voltage. Because an n-type semiconductor acts like a material with negative charges and a p-type semiconductor acts like a material with positive charges, the charge carriers of these materials will flow toward the opposite voltage polarity. With one polarity, the charge will flow across the junction, and when it is switched, the charge will not flow across the junction. 8.39 An instantaneous dipole is a small separation of charge (positive and negative) that happens because of a fluctuation in the positions of the electrons of an atom or molecule. If the electron density is momentarily in one region, that portion of the molecule will be negatively charged and the portion where the electrons “would otherwise be” will be positively charged, leading to a transient dipole.

540

Appendix K

(b) increases (c) does not change (d) increases 8.57 (a) O2 (b) SO2 (c) HF (d) GeH4 8.59 The vapor pressure will be low and the boiling point will be high. 8.61 Substance A has the highest vapor pressure. Substance C will have the next highest vapor pressure and finally, Substance B will have the lowest vapor pressure. 8.63 The liquid would look like this drawing.

8.65 For addition polymerization to occur, there must be a carbon–carbon double bond present, so this would be the thing to look for in a monomer.

(b) The temperature is roughly 35°C, so we can extrapolate from the graph and estimate the pressure to be 12,000 torr, or roughly 16 atm.

8.67 Polymerization reactions can be viewed as the growing of a molecular chain. The reactions that increase the length of the chain and leave it capable of continuing to grow are called propagation. Eventually a reaction occurs that doesn’t allow the chain to keep growing, and that is the chain termination reaction.

(c) The storage tanks will need to be strong enough to withstand high internal pressure. 8.91 The pressure from evaporation of the full amount in a closed room of this size is roughly 17 torr. This is smaller than the vapor pressure, so the entire volume will evaporate.

8.69 Because the atoms connected to the polymer backbone are all the same. They cannot be distinguished, so there are no isotactic or syndiotactic forms.

8.93 Answers will vary somewhat depending on the websites found. Many solar panels are based on p-n junctions with large areas.

8.71 There are always two products—the polymer that forms and some small molecule, often water.

8.95 (a) The alternating chain for syndiotactic polystyrene allows better packing of the polymers so that the melting point is higher. (b) Designs that require higher performance, including a higher melting point for the plastic, might require the syndiotactic polymer. One possible example would be medical devices that must be exposed to high temperatures for sterilization.

8.73 In a block copolymer, the different monomers present are in groups that are still part of the backbone. In graft copolymers, the different groupings of monomers are attached to the backbone as side chains. 8.75 Thermoplastics soften (and eventually melt) at high temperatures. 8.77 When a thermosetting polymer is heated, it forms bridges or chemical links between chains. These links are called cross-links, and they tend to prevent the resulting polymer from softening when heated. 8.79 Requires web research—answers will vary though it is unlikely that students will report that melting an eraser will allow it to be remolded and still work. 8.81 Precise answers will vary, but all should show that the initial reception to the announcement tended to be cool. Many scientists were skeptical about the existence of the spherical buckminsterfullerene, or “buckyball,” until it was made in bulk quantities. 8.83 The “heat-and-beat” method is used for solid-state reactions. The atoms of solid systems are held in place, so heat is needed for the atoms to migrate enough to react with other atoms and to form new bonds. 8.85 Web research will vary. 8.87 (a) about 350 mmHg (b) Ethanol has stronger intermolecular forces. (c) about 84°C (d) increases 7000

8.89 (a)

8.99 The magnet is providing an external field that is inducing an alignment of electrons in the pin. Polarization is similar in that an external field is used to alter the electron distribution of a molecule. 8.101 To measure the degree of polymerization, the engineer would have to measure the molecular mass of the polymer and divide by the molecular mass of the monomer. 8.103 Cryolite is a very useful flux for separating the Al and O in aluminum oxide. 8.105 Answers will vary depending on the websites found. But a typical doping level might be on the order of one dopant atom per 1,000,000 silicon atoms. To produce an n-type semiconductor, we might use phosphorous as the dopant. We will calculate our answers using these choices. (a) Doping 1 metric ton of silicon with phosphorous at a level of one part per million would require just 1.1 g of phosphorous. (b) The mole fraction of the phosphorous dopant would be 1026. Chapter 9

6000 P (torr)

8.97 To avoid creating an n-type or a p-type semiconductor, we must make sure the dopant has the same number of electrons as silicon. To lower conductivity, we must increase the size of the band gap, and so the dopant has to be smaller than silicon. The only candidate is carbon, whose atoms have four valence electrons and are smaller than silicon atoms.

9.1 There are many possible reasons why these two items could be correlated. The main idea is that the productivity of an individual worker increases, depending on the amount of energy available. A productive economy is usually a welldeveloped economy—so the energy spent per worker is related to the level of development in an economy.

5000 4000 3000 2000 1000 0 ⫺100 ⫺80 ⫺60 ⫺40 ⫺20

0

20

40

T (°C)

From the graph, we can estimate the normal boiling point to be about 245°C.

9.3 (a) Europe would probably be relatively similar. (b) The main difference between the United States and a developing country would be the total size of the graph. Chances are also good that the percentage of waste in the developing country would be larger. 9.5 The energy imports shown in Figure 9.2 are equivalent to 5.3 billion barrels of oil a year. This means we must

Appendix K

541

import nearly 15 million barrels a day. (Note that this number is high because some of our energy imports are in forms other than oil.)

9.53 (a) Two moles of the compound are formed. (b) The equation is not balanced. (c) The standard state of oxygen is O2.

9.7 The internal energy is determined by the total of the kinetic and potential energy of the atoms in the system.

9.55 –217.3 kJ/mol; 262 J is released.

9.9 6.0 3 10220 J

9.59 2126 kJ/mol

kg m1 s22 Force 9.11 P 5 ———— 5 ————— 5 kg m21 s22 so P 5 kg m2 s22, Area m2 which is an energy unit. 9.13 A hydrocarbon is a compound composed only of hydrogen and carbon. Many fuels are hydrocarbons. 9.15 (a) 3.61 3 103 cal (b) 5.9 3 105 J (c) 3.00 3 1025 Btu 9.17 16.83 3 104 kJ 9.19 A negative sign on DE means that internal energy has been lost by (or has left) the system. 9.21 Work plays a more important role in society and therefore is more valued. Consider how fuel efficiency is measured for a car. It is in miles per gallon—miles driven represent work. There is little value placed on how warm the engine has gotten. 9.23 (a) w 5 2836 J (b) w 5 230 J. So more work is done in the system in (a). 9.25 Increased efficiency does not always imply lower household energy demand. In some cases, people may choose to own more appliances such as televisions or computers. Even though each appliance is more efficient, because more are present there is no decrease in total energy demand. 9.27 According to one site, only 8% of energy use in the United States is attributable to lighting, and of that small percentage, only 27% is from residential lighting. So only slightly more than 2% of energy use in the United States can be attributed to the lighting in residences. Although it might make sense to conserve energy in this way, it’s not likely to have a major impact on total energy use. 9.29 Calorimetry is the science of measuring heat flow. It is the primary way that we determine changes in the energy of systems. 9.31 3.50 3 105 J 9.33 5.2 °C 9.35 0.52 J g–1 °C –1 9.37 383 J °C –1 9.39 700 MJ 9.41 Under constant pressure 9.43 An exothermic process releases heat to the surroundings, and an endothermic one absorbs heat from the surroundings. 9.45 Because a phase change is occurring, there is no change in temperature. Heat flowing into the substance is used to overcome intermolecular forces. 9.47 23.4 3 102 kJ/mol 9.49 5.87 kg acetone 9.51 A formation reaction is one in which exactly one mole of a chemical compound is formed from its elements in their standard states.

542

Appendix K

9.57 C2H2(g) 1 H2(g) : C2H4 DH ° 5 2175 kJ 9.61 The sum of the bond energies in the products must be greater than those in the reactants. 9.63 3.2 3 103 kJ 9.65 699 kJ 9.67 480 kJ 9.69 31.9 g/mol 9.71 Because in transportation, not only are products (and their mass) moved around, the fuel in the vehicles must also be moved around. Keeping the mass of the fuel as low as feasible means that as little energy as possible is spent moving fuel (and a greater percentage is spent moving products). 9.73 The reactions inside a battery are oxidation–reduction reactions. 9.75 These terms refer to the ability to recharge a battery. A primary battery cannot be recharged, and a secondary battery can be. Ordinary alkaline cells are primary batteries, whereas car batteries are secondary batteries. (Other examples are also possible.) 9.77 9°C 9.79 (a) Within experimental tolerances, the temperature change should be the same. Twice as much heat is released, but twice as much solution is present to absorb it. (b) In this case, the temperature rise will be greater. Twice as much heat is released, and here the solution volume is the same as in the original trial. 9.81 33.0°C 9.83 54.7 kJ is released. 9.85 106 g CH4 9.87 (a) 23°C (b) 360 cm (c) 0.038% (d) Over the length of the wire, there is an expansion of 0.139 cm. This is not likely to be significant except in the most exacting of designs. 9.89 First, we would have to assume that the heat transfer process is 100% efficient. Then we must calculate the amount of heat it takes to cool the 500 kg of fruit (using some reasonable value for the initial temperature to calculate DT) using the equation, Heat 5 mass 3 specific heat 3 DT. Then this amount of heat would go into evaporating the refrigerant, whose mass can be calculated because we know that 680 J is absorbed per gram evaporated. 9.91 Upon drawing the Lewis structures, it becomes apparent that the total number of chemical bonds for reactants and products in this reaction remains the same—that is, there are 3 B!Br and 3 B!Cl bonds in both reactants and products. Although there may be slight changes because of the specific molecular environment in which the bonds are found, the net change is going to be small.

9.93 For any reaction, the heat of the reaction is given (according to Hess’s law) by the sum of the heats of formation multiplied by their “signed stoichiometric coefficients.” In this case, the enthalpy change for the reaction has been measured. If every other heat of formation other than silicon nitride is known (or can be looked up), we can solve the Hess’s law problem for the heat of formation of silicon nitride. 9.95 This thermochemical equation refers to the energy of combustion of 2 moles of octane. To obtain the energy released per gram we must determine the molar mass of octane and convert from moles to grams. 9.97 Generally, it would be best to have a material with high heat capacity. Such a material would be capable of absorbing a large amount of heat without having a large change in temperature according to the equation, Heat 5 mass 3 specific heat 3 DT. Having the temperature stay relatively constant would have an advantage because the material would not get too hot for safe use when people might touch it. 9.99 5.25 3 1016 kJ 9.101 1.26 3 105 kJ per gallon. Roughly 69% efficient. Assumptions, such as no loss to friction, etc., are major and would generally tend to push this number toward lower efficiency. Chapter 10 10.1 From an environmental perspective, the recycling process involves the input of energy and often other resources that may have an environmental cost. 10.3 Most uses are associated with the production of fibers for things like durable carpeting or insulating fibers such as fiberfill. 10.5 Web research required—answer depends on site found. 10.7 (a) Spontaneous: this is the sublimation of dry ice. (b) Not spontaneous: table salt is solid at room temperature. (c) Not spontaneous: table salt does not dissociate at room temperature (d) Not spontaneous: CO2 is stable and does not dissociate into its elements. 10.9 Combustion of a fuel is spontaneous, for example, once it has been initiated by a spark. Spontaneous does not imply instantaneous upon mixing. 10.11 (a) Nonspontaneous. Opposite: Formation of O2 from oxygen atoms. (b) Nonspontaneous. Opposite: A tray of ice melts on a warm day. (c) Nonspontaneous. Opposite: An acid and base form a salt solution after the reaction is complete. (d) Spontaneous. (e) Nonspontaneous. Opposite: SO3 and water react to form sulfuric acid. 10.13 (a) T h i s p r o c e s s i s n o t s p o n t a n e o u s a t h i g h temperatures. (b) This process is not spontaneous at very high temperatures. (c) This process is always spontaneous. (d) This process may not be spontaneous at high temperatures.

10.15 No, at the very least the system must include oxygen (often from air). 10.17 Alice said, “With a wall so narrow your balance must sometime fail you—chances are you will soon fall. The ground is not so narrow, so you cannot fall from it. And as for the king—he may mean well but I doubt he is a magician. If you break into 1000 pieces what will become of you if the king finds only 995 of them? I don’t think you should sit on that wall waiting to find out.” 10.19 He found that when he calculated the heat around the whole path divided by temperature (q/T ), the result was zero. 10.21 Nature spontaneously tends toward a more probable state. 10.23 (a) CO2 vapor has higher entropy because gases are more randomly arranged than solids and thereby have more microstates. (b) Sugar dissolved in water will have higher entropy because the molecules are dispersed, and more dispersed systems have more microstates. (c) The mixed liquids have more entropy because they have dispersed into each other. 10.25 DS , 0. This process is the opposite of mixing, so the sign on DS must be negative. 10.27 In each case the side of the equation with more moles of gas will have the higher entropy. So the signs would be (a) positive, (b) positive, (c) negative, (d) negative, and (e) positive. 10.29 The reduction of entropy in the leaf itself as it grows must be accompanied by a larger increase in entropy elsewhere. Decomposition of the leaf after it falls from the tree leads to a large increase in entropy. 10.31 The thermodynamic entropy increase occurs mainly in heating the surroundings. 10.33 Once in a deck, playing cards would not become a disarrayed pile, whereas atoms and molecules will constantly move and mix, leading to the more probable state without any intervention. 10.35 It increases. 10.37 To convert heat into work requires a decrease in randomness, which would lower entropy. 10.39 The exothermic nature of the change means that heat is released to the surroundings. This leads to an entropy increase in the surroundings which is greater than the decrease in the entropy of the system. 10.41 It generally has no influence on the DS of the system. The exothermic part has a major effect on the entropy of the surroundings. 10.43 The reaction, Br 2(g) 1 Cl 2(g) : 2 BrCl(g), has the greater entropy change. There are more possible “pairings” for the heteronuclear molecules, so in each case those have the higher entropy DS°. 10.45 The process is endothermic. 10.47 (a) DS° 5 2152.4 J/K (b) DS° 5 2268.0 J/K (c) DS° 5 2185.0 J/K (d) DS° 5 2116 J/K (e) DS° 5 1134.2 J/K 10.49 DS° 5 2114.6 J/K. This negative value for entropy change arises from the fact that water molecules solvate ions in solution, thereby creating order.

Appendix K

543

10.51 (a) DS° 5 164.0 J/K (b) Plants carry out this type of reaction in conjunction with other reactions that increases entropy. As long as the overall entropy change of the universe is negative, the coupled reactions can proceed spontaneously. 10.53 H2C2O4(s) , C2H5OH(,) , CH4(g). The entropy of a gas is generally much greater than that of a liquid, which in turn is generally greater than that of a solid. 10.55 (a) At the molecular level, the water molecules in the solid (ice) are becoming less ordered as it melts. The ice water is still at 0°C because all of the heat that flows into it melts the ice. (b) Inside the container, the entropy must be increasing because the system is isolated. 10.57 When the process takes place at constant pressure and temperature. 10.59 (a) 514 kJ (b) 21295 kJ

(c) 35 kJ

10.61 Both the enthalpy change and the entropy change are positive. At lower temperatures, the enthalpy term dominates. But at higher temperatures, the TDS term becomes more important, leading to a negative DG. 10.63 A reversible change is a small change where the system can be restored to its original state. An irreversible change places the system in a state from which the original state cannot be restored. 10.65 (a) DS° 5 297.4 J/K (b) DG° 5 21226.5 kJ (c) At high enough temperatures, the TDS term could dominate this equation. (d) Using the equation DG° 5 DH° 2TDS° might be questionable for determining this range because it requires that DH ? f (T). 10.67 (a) SiH4(g) 1 2 O2(g) : SiO2(s) 1 2 H2O(g) (b) DS° 5 2195.3 J/K (c) DG° 5 21370.8 kJ, so the process is spontaneous. (d) For methane combustion, DG° 5 2800.9 kJ which is not nearly as negative as for silane. Both reactions are enthalpy driven. 10.69 (a) DG° 5 111.7 kJ (b) DG° 5 18.3 kJ (c) DG° 5 2817.6 kJ (d) DG° 5 240.42 kJ (e) DG° 5 1617.0 kJ 10.71 For dissolving, DG°(NaCl) 5 29 kJ and DG°(AgCl) 5 55.7 kJ. The negative value of DG° for NaCl suggests that it dissolves, whereas the positive value of DG° for AgCl suggests that it doesn’t. 10.73 (a) positive (b) negative (c) zero (d) positive (e) negative 10.75 (a) spontaneous at lower temperature (DH , 0, DS , 0) (b) spontaneous at higher temperature (DH . 0, DS . 0) (c) not spontaneous at any temperature (DH . 0, DS , 0) (d) spontaneous at higher temperature (DH . 0, DS . 0) 10.77 No. An endothermic reaction cannot be spontaneous if it also has DS , 0. 10.79 If the long chain breaks into a shorter chain and a small piece, randomness has increased. 10.81 DS 5 84.4 J/K/mol 10.83 Tbp 5 2153°C. 10.85 (a) DG° 5 16.2 kJ (b) DH° 5 17 kJ

544

Appendix K

10.87 (a) False. An exothermic reaction is often, but not always, spontaneous. (b) False. When DG° is positive, the reaction cannot occur unless it is driven by an external system. (c) False. DS° is positive for a reaction in which there is an increase in the number of moles of gas in going from reactants to products. (d) False. If DH° and DS° are both negative, DG° will be negative at low temperatures. 10.89 (a) If we assume an initial temperature of 25°C, then the maximum temperature will be ,48°C. (b) Probably not, because heat will be dissipated to the surroundings and whatever products arise from the exothermic process will also have some heat capacity. (c) Assuming the process occurs at constant pressure, DH is –1500 J. (d) DS 5 q/T 5 1500/298 5 5.0 J/K 10.91 (a) 467.9 kJ (b) 87.4 J/K (c) 301.2 kJ (d) T . 838.6 K 10.93 Use the sample to carry out a combustion reaction in a calorimeter. From the heat of combustion, along with heats of formation of the products (carbon dioxide and water), you can calculate the heat of formation of cyclohexane. 10.95 The equation DG° 5 DH° 2 TDS° provides a relationship between free energy, enthalpy, and entropy. If we can obtain values for any two of these variables (in this case the entropy and free energy changes) for a given reaction, we can calculate the third. So if we calculate DS° for the formation reaction of the substance in question, we can use that value along with the known value of DGf° to find DHf°. (We would use the standard temperature of 298 K.) 10.97 The key feature lies in the sign of the enthalpy and entropy change. If they have the same signs (both positive or both negative), the spontaneity of the process will be temperature dependent. Melting snow is an endothermic process (DH . 0), but it increases entropy (DS . 0), so it will tend to occur at warmer temperatures. Snow formation is exothermic but lowers entropy, so it will tend to occur at lower temperatures. 10.99 (a) 2 F2 1 2 H2O : 4 HF 1 O2 (b) DG° 5 2618 kJ, so the reaction is spontaneous. (c) DH° 5 2512 kJ (d) So, 233 kJ of heat is released when this much F2 is bubbled through. 10.101 (a) MnO: manganese(II) oxide; MnO2: manganese(IV) oxide; Mn2O3: manganese(III) oxide (b) 3 MnO 1 1/2 O2 : Mn3O4; 3 MnO2 : Mn3O4 1 O2; 3 Mn2O3 : 2 Mn3O4 1 1/2 O2 (c) for the MnO reaction, DG° 5 2194 kJ; for the MnO2 reaction, DG° 5 1113 kJ; and for the Mn2O3 reaction, DG° 5 178 kJ. These numbers suggest that MnO2 is the most stable form. Chapter 11 11.1 An allotrope is a single physical and chemical form of an element. Both diatomic oxygen and ozone are allotropes of the element oxygen. 11.3 Ozone is considered a pollutant when it forms near the surface of the earth in the troposphere. When it forms in

the stratosphere as part of the Chapman cycle, however, it has the beneficial attribute of protecting the planet from harmful UV radiation.

11.37 Looking at the figure of the cube, 2 2

11.5 There is no net reaction—and therein lies the reason it is called a cycle.

2

11.7 The breaking of the double bond in diatomic oxygen requires the most energy and the shortest wavelength light.

2

11.9 No—it is a seasonal phenomenon. The depletion of stratospheric ozone that is referred to as the ozone hole occurs in the springtime over polar regions.

2

we can calculate the surface area of the whole cube as 6 3 (2 3 2) 5 24 cm2. The surface area of each cube in the group of eight smaller cubes is 6 3 (1 3 1) 5 6 cm2, and for eight such cubes, the surface area is 8 3 6 5 48 cm2. Dividing the eight small cubes into even smaller cubes will increase the surface area again. Consider a log versus all the toothpicks that could be made from it. The toothpicks are analogous to the smaller cubes in this picture. They would have large surface area compared to the log, and it would be much easier to start a fire with the toothpicks than with a complete log.

11.13 An instantaneous rate is measured at one moment in a reaction, whereas an average rate is measured over some finite period of time. The hot dog eating rate of 5 per minute over the duration of the contest is an example of an average rate. The reported rates of 6 hot dogs per minute at 1.0 minute and 3 hot dogs per minute at 3.0 minutes are examples of instantaneous rates at those points in the contest.

11.17 The rate of oxygen consumption is 5.29 3 1024 mol L21 s21. NO is produced at the rate of 4.23 3 1024 mol L21 s21. Water is produced at the rate of 6.35 3 1024 mol L21 s21. 11.19 For the first 8.3 minutes, the rate is 0.51 L/min. Over the entire 19-minute experiment, the rate is 0.605 L/min. 11.21 Estimating the slope from a graph, we obtain a rate of 0.37 torr/min. 11.23 The rate will increase by a factor of 9 if [A] is tripled. The rate will decrease by a factor of 4 if [A] is halved. 11.25 (a) 1st order in X, 1st order in Y, and 1st order in Z; 3rd order overall (b) 2nd order in X, 0.5 order in Y, and 1st order in Z; 3.5 order overall

11.39 [N2O5] 5 0.146 M 11.41 0.4 mg 11.43 260 mg 11.45 (a) 0.16 mol/L (b) 90 s (c) 120 s 1 1 11.47 We have ——— 5 kt1/2 and t1/2 5 ——— The half-life k[ X ]0 [ X ]0 is concentration dependent and will change during the course of the reaction. 11.49 Looking at the various possible plots, we find a straight line for the plot of ln[picloram] vs. time so the reaction is first order. k 5 1.7 3 1022 days21, and t1/2 5 42 days. 11.51 This problem requires plotting data to find straight lines. Test plots for first- and second-order kinetics are as follows: –6

(c) 1.5 order in X and 21 order in Y; 0.5 order overall

–6.2

(d) 1st order in X and 22 order in Y; 21 order overall

11.29 x 5 1 and y 5 2

–6.4 –6.6 ln P

11.27 Solving the rate law for k, we have k 5 rate/[H2][Br2]. Inserting the given units for the rate and mol/L for each of the two concentrations gives us units of L mol21 s21 for k.

–6.8 –7

11.31 First order with respect to A, second order with respect to B, zero order with respect to C, and third order overall. k[NO22][NH41]. The 24 21 21

11.33 Rate 5 is 3.7 3 10

1 2

11.11 (a) beats per minute (b) feet per second (c) rotations per second (d) volume per millisecond.

11.15 The combustion of hydrocarbon fuels produces carbon dioxide and water, so the production of these chemicals could also be measured to determine the combustion rate of the candle.

1

1

L mol

value of the rate constant

s .

2

11.35 Rate 5 k[NO][H 2] , and the rate constant is 23 L mol22 s21.

2

–7.2 –7.4 –7.6

0

10

20

30 40 Time (min)

50

60

Appendix K

70

545

2000

11.71 (a) NO2(g) 1 NO2(g) : NO(g) 1 NO3(g)

1800

NO3(g) 1 CO(g) : NO2(g) 1 CO2(g)

1600

NO2(g) 1 CO(g) : NO(g) 1 CO2(g)

1/P (torr–1)

1400

(b) Each step is bimolecular.

1200 1000

(c) Rate 5 k[NO2]2 (d) NO3(g) is an intermediate.

800 600

11.73 To the extent that the success of the sports team parallels the progress of a chemical reaction, the reserve that enhances the success of the sports team works the same way as a chemical that enhances the rate of a reaction.

400 200 0

0

10

20

30 40 Time (min)

50

60

70

From these graphs, we can see that the plot of ln P versus time is linear, so the reaction is first order. The slope of that graph is equal to 2k. A linear regression fit to the line gives a slope of 20.0217 min21, so the rate constant is k 5 0.0217 min21. 11.53 (a) The reaction is first order. (b) k 5 3.3 3 1023 s21. (c) Initial rate 5 3.3 3 1023 M s21. 11.55 k 5 5 3 10270 s21 11.57 Ea 5 120 kJ/mol 11.59 Ea 5 3.2 3 105 J/mol 5 320 kJ/mol 11.61 This experiment would need some measurement to determine if the milk has soured—perhaps the smell of the milk could be noted at a given time during the day. The next step would be holding the temperature steady at several values. Then, assuming the only source of lactic acid is the reaction in the bacteria, the activation energy could be estimated using the Arrhenius equation. 11.63 A reaction mechanism cannot be proven correct. Even if a particular intermediate is observed, there is always the possibility that another mechanism could be proposed that also generates that intermediate. A mechanism can be much more readily proven incorrect. For example, any measurement that demonstrates the existence of an intermediate that is not part of the proposed mechanism proves that the mechanism is wrong. 11.65 Because the Chapman cycle includes species that are intermediates in the overall reaction, it must be a mechanism of the overall reaction. The molecularity of the three steps in order are O2 : 2 O (unimolecular—light has been left out of this equation) O2 1 O : O3 (bimolecular) O 3 : O 2 1 O (unimolecular—light left out in this equation) O3 1 O : 2 O2 (bimolecular) 11.67 No, this equation is clearly not a mechanism. If it were, it would imply a six-body collision. 11.69 (a) 2 NO 1 Br2 : 2 NOBr (b) the first step (c) NOBr2 is the intermediate. (d) Both are bimolecular. (e) Equation #1: Rate 5 k[NO][Br2], and Equation #2: Rate 5 k[NOBr2][NO].

546

Appendix K

11.75 Homogeneous catalysis takes place completely within a single phase, whereas heterogeneous catalysis takes place in two phases, often on the surface of a solid. 11.77 It would do the same thing—it would lower the activation energy of both the forward and reverse reactions. 11.79 In one case, the pollutant needs to be reduced (NO2), and in the other case, the pollutant needs to be oxidized (CO). It’s unlikely that a single catalyst would work well for two very different reactions such as these. 11.81 The key initiating reaction for smog formation requires light. So this is a photochemical reaction. 11.83 VOC is an acronym for volatile organic chemical. VOCs are significant in photochemical smog formation because they destroy radicals that form in some of the initiating reactions. If these radicals were not destroyed, the reactions that formed them would tend to reverse, and the smog formation would be reversed almost as soon as it started. 11.85 The presence of nitrogen oxides and VOCs are the key factors. In places where there are many automobiles, such as urban centers, there are ample sources for these two compounds. 11.87 Because several reactions are required from the initial generation of the pollutants, such as nitrogen oxides, before ozone is formed. It takes some time for these reactions to take place—so there is a lag in time before ozone forms. 11.89 Answer varies depending on location. The Environmental Protection Agency website at www.epa.gov represents a good source to answer this question. 11.91 C6H14 1 HO? : C6H13? 1 H2O C6H13? 1 O2 1 M : C6H13O2? 1 M C6H13O2? 1 NO : C6H13O? 1 NO2 The NO2 can then be recycled into the ozone generation process by creating O. In addition, by consuming NO, the reaction prevents some of the destruction of ozone as it is created (because NO reacts with ozone, as noted in Equation 11.10.c). 11.93 (a) F

Xe

F It’s a linear molecule.

(b) In this case, the partial pressures can be determined from the volumes, so the partial pressure of Xe is 459 mm Hg, and the partial pressure of F2 is 301 mm Hg. (c) nF2 5 0.0107 mol and nXe 5 0.0164 mol, so the fluorine is the limiting reactant. There are 0.0107 mol of XeF2 formed, so the theoretical yield 5 1.81 g XeF2.

(d) Precipitation reactions are usually very fast com-pared to this reaction. The need to dissociate the F2 molecules in the gas phase reaction is one reason for this. 11.95 The value observed is 0.13 ppm. Assuming atmospheric pressure of 1 atm and a temperature of 298 K, we can calculate that there are 3.2 3 1015 molecules of O3. 11.97 (a) Chances are that it slows as the process continues. (b) Eventually, the rate is the same in both directions. 11.99 The rate of ammonia formation is much greater at high temperatures. The reaction must have a high activation energy such that high temperatures are required to produce collisions with enough energy to overcome the activation barrier. This is consistent with the fact that the triple bond in N2 is extremely strong. 11.101 The reaction should be 2 Be 1 Si : Be2Si, and 401 g of Be are needed. 11.103 76 kJ/mol 11.105 First, it’s important to realize that radioactivity is essentially a first-order kinetics problem. Knowing the percentage of counts as a fraction of that seen for living organisms allows us to set up the ratio, I/Io, which is equal to e2kt. To use this relationship, we need to know the value of k or of the half-life, and then we can solve for time, t, which will be the age of the artifact. 11.107 To solve this problem, we need to consider factors that change the rate of reaction. In some cases, the presence of products will slow the reaction. Another strategy would be to lower the concentration of the chemical or chemicals in the rate law of the reaction or process that is being used. 11.109 Two things. First, the bonds fluorine atoms form with other elements are stronger than F!F bonds, making the reactions highly exothermic. Second, the activation energies for the reactions are probably not very large. 11.111 Radicals are very reactive—so for them to be useful, they must be generated in the reaction mixture (it’s nearly impossible to have a bottle of them). The high temperatures allow bonds to break and radicals to form. Chapter 12 12.1 The first reaction is the decomposition of calcium carbonate, CaCO3 : CaO 1 CO2, and this reaction is essentially the opposite reaction that takes place in the carbonation of concrete. 12.3 Different sources estimate this number differently, but roughly 5% of annual CO2 production can be attributed to concrete production. This is largely due to the decomposition of calcium carbonate. 12.5 The reaction is slow, so the activation energy for it to occur must be relatively high. 12.7 The parts of the concrete that were exposed to air are likely to be carbonated and therefore show no color change upon exposure to phenolphthalein. The parts that had been internal were not exposed to air, so the concrete should retain its original chemical composition and turn pink when exposed to phenolphthalein.

12.9 (a) In the first few minutes, orange fish will tend to swim toward the right, and green fish will tend to swim to the left. (b) After several hours, the distribution should be uniform throughout. (c) This is analogous to a dynamic equilibrium because fish swim in all directions, but there is no net change in the distribution of the locations of the different colored fish. 12.11 RateH21I2 : 2HI 5 rate2HI : H2 1 I2 12.13 The top line is the concentration of A, and the bottom line is the concentration of B. At (a), only A is present, so the rate of disappearance of A is large, as is the rate of appearance of B. At (b), the rate of the forward reaction has slowed somewhat, and the rate of the reverse reaction is becoming appreciable. Finally, at (c), the rates of the forward and reverse reactions are the same, so there is no net change in the concentration of either chemical. 12.15 (a) This inference can be made because the equilibrium symbol implies this fact. (b) This is not implied by equilibrium. Although concentrations do not change at equilibrium, this does not imply that they are all the same. (c) This is not implied either. Chemical equilibrium can be approached from either “direction”—starting with all reactants or with all products are just two possible starting points. [NO]2[Br2] [H2O]2[Cl2]2 —————— 12.17 (a) K 5 —————— (b) K 5 [NOBr]2 [HCl]4[O2] [SO3] [CO2][H2O]2 (d) K 5 ——————— (c) K 5 —————— 1/2 [CH4][O2]2 [SO2][O2] 2 3 [CO2] [H2O] (e) K 5 ———————— [C2H5OH][O2]3 12.19 (a) K5[Ca21][CO322] (b) K 5 [Ag1][Cl2] [Zn21] (c) K5[Mg21]3[PO432]2 (d) K 5 ———— [Cu21] 12.21 (a) Reactants are favored. (b) Products are favored. (c) Reactants and products are present in roughly in equal concentrations. 12.23 (a) Lead is more likely to precipitate sulfate ions. [O2] 12.25 (a) For the reverse reaction, K 5 ———— [CO]2 1 (b) For the reverse reaction, K 5 —————— [Ag1][Cl2] [HCO32] [H2CO3] 12.27 (a) K1 5 ——————— and K2 5 ——————— 22 1 [H1][HCO32] [H ][CO3 ] (b) Summing the two reactions leads to CO322 1 2 H1 N H2CO3 [H2CO3] (c) The equilibrium expression, is K3 5 ——————— , [H1]2[CO322] which is the product of K1 and K2. [CO2] 12.29 K 5 —————— [CH4] [O2]2 12.31 Because K2 is greater than K1, reaction #2 (from problem 12.30) has a greater tendency to go toward completion. 12.33 K 5 1.2 12.35 K 5 0.080

Appendix K

547

12.37 The equilibrium concentrations are [O2] 5 0.100 M and [O] 5 1.7 3 10220 M. 12.39 [NO] 5 0.0020 M and [N2] 5 [O2] 5 0.099 M. 12.41 [NO] 5 0.002 M, [N2] 5 0.149 M, and [O2] 5 0.0491 M. 12.43 [NO] 5 0.0017 M, [N2] 5 0.029 M, and [O2] 5 0.029 M. 12.45 [H2CO3] 5 0.0998 M, [H1] 5 0.00021 M, and [HCO32] 5 0.00021 M 12.47 [Ca21] 5 0.0 M and [CO322] 5 0.010 M. (Within the significant figures we have available, the concentration of calcium ion is zero.) 12.49 (a) Shifts the equilibrium toward products. (b) Shifts equilibrium toward products. (Adding water will reduce [H3O1] and [F2], making Q , K.) (c) More products will be formed. (d) The equilibrium shift will be toward products. (e) The equilibrium shift will be toward products. (f ) Addition of a catalyst has no effect on the equilibrium. 12.51 (a) Shift toward reactants. (b) Shift toward the reactant. (c) There is no gas involved in this equilibrium, so pressure changes do not affect it directly. 12.53 Because the process is endothermic, increasing the temperature will drive the reaction to the right. Adding more NH4HS to the flask will also drive the reaction to the right. If additional NH3 is added to the flask, the reaction will be driven to the left. If H2S is removed from the flask, the reaction will be driven to the right, and the pressure of NH3 will increase. 12.55 (a) No effect on equilibrium concentrations. (b) Shifts toward products. (c) Shifts toward products. (d) Does not affect the equilibrium. 1

2

1

2

12.57 (a) AgBr(s) N Ag (aq) 1 Br (aq); Ksp 5 [Ag ][Br ] 21

2

(b) Cu(OH)2(s) N Cu (aq) 1 2 OH (aq); Ksp 5 [Cu21][OH2]2 (c) BaSO4(s) N Ba21(aq) 1 SO422(aq); Ksp 5 [Ba21][SO422] (d) PbCrO4(s) N Pb21(aq) 1 CrO422(aq); Ksp 5 [Pb21][CrO422] (e) Mg3(PO4)2(s) N 3 Mg21(aq) 1 2 PO432(aq); Ksp 5 [Mg21]3[PO432]2 12.59 1.4 ppm

(a) HCN & CN2 and H2O & OH2 (b) H3O1 & H2O and H2CO3 & HCO32 (c) H2S & HS2 and CH3COOH & CH3COO2 12.75 (a) CH3COOH(aq) 1 H2O(,) N H3O1(aq) 1 CH3COO2(aq) [H3O1][CH3COO2] Ka 5 —————————— [CH3COOH] (b) C2H5COOH(aq) 1 H2O(,) N H3O1(aq) 1 C2H5COO2(aq) [H3O1][C2H5COO2] Ka 5 —————————— [C2H5COOH] (c) HF(aq) 1 H2O(,) N H3O1(aq) 1 F2(aq) [H3O1][F2] Ka 5 —————— [HF] (d) HClO(aq) 1 H2O(,) N H3O1(aq) 1 ClO2(aq) [H3O1][ClO2] Ka 5 ——————— [HClO] (e) H2CO3(aq) 1 H2O(,) N H3O1(aq) 1 HCO32(aq) [H3O1][HCO32] Ka 5 ———————— [H2CO3] 12.77 [H3O1] 5 0.008 M 12.79 pH 5 2.95, 1.1% ionized 12.81 pH 5 2.61 12.83 pH 5 9.93 12.85 (a) Because the value of K is less than 1, it suggests that equilibrium lies toward reactants, and therefore DG° should be a positive number. (b) Because we want the standard free energy change, we use T 5 298 K, and DG° 5 12.1 kJ mol21. 12.87 This was probably an error in units. The value for R is in J and the DG° value is in kJ. If the numbers obtained for DG° are converted to J, the differences among the reactions become more obvious. The correct K values are 5.8 3 108, 1.8 3 105, and 57, respectively. 12.89 [CO2] 5 2.5 3 1023 M

–4

–7

–7

12.61 (a) 2 3 10 M (b) 6 3 10 M (c) 6 3 10 M 12.63 (a) K sp 5 3.0 3 10 H2O 12.65 (a) 4.9 3 10211 244

(c) 2.8 3 10

253

(b) 1.3 3 10

225

g HgS /100 g

(b) 1.4 3 10216 (d) 1.6 3 1025

12.67 BaSO 4 will precipitate first. A saturated BaSO 4 solution has a Ba21 concentration of 1.05 3 10–5 M, which will be the concentration of Ba21 when the lead starts to precipitate. This means that 99.99% of the Ba21 has precipitated. 12.69 Web question—answer will vary depending on site found. The key idea is that calcium carbonate is less soluble in water at higher temperatures, leading to the formation of scale in equipment used to boil or heat water. 12.71 (a) NO32 (b) OH2 (c) SO422 (d) HCO32 (e) H2O

548

12.73 For the purpose of this guide, we’ll just list them rather than drawing brackets. (Conjugate acid written first in each case.)

Appendix K

12.91 (a) H2S(aq) 1 Fe(NO3)2(aq) N FeS(s) 1 2 HNO3(aq) [HNO3]2 (b) K 5 ———————— [H2S][Fe(NO3)2] (c) Removal of one or both of the products will prevent it from reaching equilibrium. 12.93 (a) Proceeds toward products. (b) Proceeds toward products. (c) Proceeds toward reactants. (d) Shifts equilibrium toward products. (e) Shifts equilibrium toward products. (f ) Shifts equilibrium toward products. 12.95 An increase in pressure increases the collision frequency of the gaseous molecules. Because 2 NO2 molecules react to form 1 N2O4, an increase in the number of collisions will tend to increase the probability that this reaction will occur until the collision frequency returns to its original value.

12.97 8.4 3 1026 g Al2(SO4)3 per gallon 23

12.99 Molar solubility 5 3.29 3 10

27

M and Ksp 5 1.42 3 10 .

between the amalgam in the filling and the exposed aluminum occurs—producing a noticeable electric current.

12.101 Measure out the same small mass of the three white solids. Add distilled water dropwise to each and count the number of drops it takes to dissolve each of them. The one that is most soluble will take the fewest drops; the one that is least soluble will take the most drops.

13.19 There must be a conductive metal present for the electrons to move in a circuit and provide the electrons for the oxidation–reduction reaction.

12.103 If you have a pH meter, by reading the pH of the solution you would be able to do an equilibrium calculation to determine the initial concentration (because you would be able to look up the Ka of this acid.)

13.23 A chemical that appears at the bottom of the standard reduction potential table with a sizeable negative value means that the chemical being considered has a strong tendency to be oxidized (and very little tendency to be reduced). This observation is in line with expectations for alkali metals that tend to lose an electron easily to form mono-positive cations.

12.105 (a) [PCl 3] 5 0.24 M, [PBr 3] 5 1.24 M, [PCl 2Br] 5 [PClBr2] 5 0.76 M. (b) This reaction is entropy driven. 12.107 Kc 5 [H2O(g)] 5 0.0159 Chapter 13 13.1 Normally, the first location where rust shows is toward the bottom of the car, particularly in the wheel-wells. These areas of the car encounter water (with salt as an electrolyte where snow removal is required), and the presence of water is necessary for the electrochemical reactions associated with corrosion to occur. 13.3 Examples will depend on Internet research. 13.5 (a) The oxidation half-reaction is Cu(s) : Cu21(aq) 1 2 e2, and the reduction half-reaction is Ni21(aq) 1 2 e2 : Ni(s). (b) The oxidation half-reaction is Ba(s) : Ba21(aq) 1 2 e2, and the reduction half-reaction is Fe31(aq) 1 3e2 : Fe(s). 13.7 It provides a source of charge so that the electrical circuit can be completed. 13.9 The anode is the site of oxidation. Generally, the site of oxidation in a galvanic cell will release positive ions into solution. So to maintain overall charge neutrality in the solution, nitrate ions from the salt bridge would be released into this compartment. 13.11 (a) Ni(s)&Ni21(aq)(0.20 M)'Ag1(aq)(0.50 M)&Ag(s) (b) Cu(s)&Cu 21(aq)(0.20 M)'PtCl 622(aq)(0.10 M), PtCl422(aq)(0.10 M), Cl2(aq) (0.40 M)&Pt(s) (c) Pb(s)&PbSO4(s)&SO422(aq)(0.30 M)'Cl2(aq) (0.20 M)&AgCl(s)&Ag(s) (d) Zn(s)&Zn 21(aq)(0.125 M)'H 1(aq)(0.010 M)& H2(g)(1.0 atm)&Pt(s) 13.13 The anions in this case are spectator ions, so they don’t need to be included in the cell notation. 13.15 The electromotive potential in galvanic corrosion is supplied largely by the electrical contact of two metals with different electromotive activities. In uniform corrosion, the same metal (normally iron) undergoes both oxidation and reduction, so only one metal is required for it to happen. 13.17 The aluminum wrapper has an oxide coating on the aluminum, and aluminum oxide is not reactive. When a person bites the aluminum, however, the much more reactive aluminum is exposed, and an electrochemical reaction

13.21 As long as the SHE was still the standard, all values would shift by 3.00 V.

13.25 (a) 1.33 V (b) 0.02 V (c) 0.87 V 13.27 1.043 V 13.29 (a) E 5 1.05 V (b) E 5 0.39 V (c) E 5 0.604 V 13.31 (a) Sn21(aq) 1 2 e2 : Sn(s) and Fe(s) : Fe21(aq) 1 2 e2 (b) Iron will be the anode and tin will be the cathode. (c) E 5 0.32 V 13.33 The memo would need to point out the chemical properties (such as corrosion resistance) of gold that make its use important. The physical properties of gold that are useful in microchip applications (electrical resistance, for example) could also be included. 13.35 Because corrosion is an example of an electrochemical reaction, an electrolyte in concrete assists the corrosion process. Chloride-containing salts such as CaCl2 provide chloride ions as an electrolyte. 13.37 Use the equation, DG 5 nFE. 13.39 (a) 2237 kJ (b) 253.0 kJ (c) 22.190 3 103 kJ 13.41 DGf° (Pb21) 5 225 kJ mol21 13.43 (a) spontaneous (b) not spontaneous (c) not spontaneous (d) spontaneous 13.45 (1) K 5 4.3 3 1045, (2) K 5 1.0 3 1025, (3) 4.3 3 1020 Looking at these values, we can see that the (1) 5 (2) 3 (3). This result suggests that if a calculator cannot handle the inverse log calculation, one need only break it into two smaller calculations and multiply the results. 13.47 Ksp 5 1.7 3 10210. This value is quite close to the tabulated value of 1.8 3 10210. 13.49 (a) K 5 3 3 106, so this reaction could be observed. (b) K 5 2 3 10 216 , so this reaction would not be-observed. 13.51 A cell based on the reaction, 3 Ag1(aq) 1 Cr(s) : 3 Ag(s) 1 Cr31(aq) has E° 5 1.54 V, so it is a good candidate. One set of concentrations that would work is to let the silver ion concentration be 0.50 M and the concentration of chromium(III) be 1.2 M. The sketch should reflect a calculation using the Nernst equation. NOTE: Other combinations are certainly possible.

Appendix K

549

13.53 Because the acidic electrolyte used in older batteries is replaced by a basic (or alkaline) electrolyte. 13.55 The electrolyte in alkaline batteries includes MnO2. In acidic solution, the half-reaction MnO2(s) 1 4 H1(aq) 1 2 e2 : Mn21(aq) 1 2 H2O has a positive standard reduction potential of 1.23 V. So this reaction would tend to proceed when coupled with other half-reactions. Also, because the half-reaction for Zn21&Zn has a negative standard cell potential, Zn will tend to be oxidized in acidic solution. 13.57 Batteries are used to do work. If the battery itself is massive, the amount of work needed to move it (along with whatever else is part of the device) would be needlessly high. Moreover, in consumer products, lightweight devices (to allow them to be easily carried) are also important. 13.59 PbSO4(s) forms as a product. This material can be converted back to either Pb(s) or PbO2(s) as long as it remains at the electrode. If the battery is jarred, however, and the PbSO4 is removed from the electrode mesh, it can’t be converted back and the battery slowly degrades. 13.61 (a) The paragraph should emphasize that battery (a) has a higher voltage when fully charged (which could be beneficial). The long lifetime of battery (d) would be the most salient factor to be noted. (b) The zinc electrode is the component that limits the lifetime of the battery (because oxygen is derived from the air.) Battery (d) must have a larger surface area for the zinc electrode than that in battery (b). 13.63 Requires Internet research, and answers will vary depending on the site found. 13.65 Because oxidation–reduction reactions are already moving electrons between various reactants to form products, it is relatively easy to force electrons to move in the opposite direction by applying an external electrical potential. By contrast, in acid–base chemistry, the water formed as a product is very stable, and breaking the strong O!H bonds in the water molecule is difficult. Thus, it is not easy to reverse this type of reaction. 13.67 It is passive electrolysis. The quickest way to see this is that the electrodes are carbon—they merely provide a way for electrons to move in the system (though oxygen produced in the process reacts with the carbon and slowly erodes it). 13.69 Because electrical energy is being forced into the electrolysis cell via an external potential, the cell reaction can have zero potential. A galvanic cell with zero potential is not possible because there would be no driving force for reaction. 13.71 Rotating the barrel shifts the small parts and therefore prevents individual parts from being “plated” for longer times than other individual parts. This would lead to uneven coatings. 13.73 Depends on Internet research. Some electroplating companies do emphasize cosmetic silver plating, so the paper could support the assertion. It is more likely that the site found will have more industrial applications for silver plating, and a paper would reflect that fact.

550

Appendix K

13.75 1.1 mol e2 13.77 0.36 kWh at 0.30 V; 1.1 kWh at 0.90 V 13.79 37 g Cu 13.81 1.772 kg Sn 13.83 18 kWh 13.85 34 PbSO4 13.87 Mass of coating 5 4.0 3 1024 g; plating time 5 0.87 s 13.89 The key to corrosion prevention is that the paint must form a strong layer that bonds well to the metal. Additives are chosen to achieve this. Some paints are also designed to minimize the spread of corrosion. These include additives that prevent the flaking of iron(III) oxide from the iron. If the rust doesn’t flake off, additional corrosion can be avoided. 13.91 The main issue would be whether or not any additives could be found that could counteract the tendency of the corrosion prevention agent to make paint brittle. Brittle paint will crack and peel under weather conditions, so the additive will not be helpful if the paint has cracked and fallen off the painted surface. 13.93 No. Steel is mostly iron whose standard reduction potential is more negative than that of nickel, so it would still be the iron that is oxidized if it is coupled with nickel. 13.95 The steel would be susceptible to galvanic corrosion because copper has a greater positive value for its standard reduction potential. This would mean that the iron in the steel would tend to corrode spontaneously when in contact with the relatively good oxidizing agent, copper. 13.97 To do this problem, you need the Nernst equation. This means we must determine the number of electrons transferred and we must be able to calculate the standard cell potential from tabulated data. We can plug this information into the Nernst equation and adjust the concentration to obtain the desired voltage. 13.99 If you can build a galvanic cell with the tin as one halfcell and the N2O as the other half-cell, the measurement of the standard cell potential would provide the best means to determine the equilibrium constant. You could also calculate this standard cell potential if the necessary reduction potentials for the relevant half-reactions are available. 13.101 The key variable is the thickness of the gold coating. A thinner coating will cost less. Other possibilities include using low voltage electroplating to reduce production costs and using barrel plating to increase the throughput of part production to lower production costs. The main pieces of information that you need to look up to make estimates of the electrolytic process are the standard reduction potentials of the half-reactions involved. 13.103 Oxygen is produced from the water solvent, 2 H2O : O2 1 4 H1 1 4 e2. The volume produced is 2.9 L. 13.105 (a) 0.30 g Ag (b) 1.7 3 1021 atoms (c) Z106 layers of atoms

Chapter 14 14.1 Cosmic rays are composed of particles (which are discrete) rather than electromagnetic radiation (which is continuous).

14.37 Loss of radon from the sample would increase the U/Pb ratio, so the measurement would underestimate the age of the rock. 14.39 1.01585 3 1010 kJ/mol

14.3 Answer may vary based on site found and how many different particles the site lists.

14.41 6.2877 3 10212 J/nucleus, or 3.7865 3 109 kJ/mol

14.5 The primary means is through photosynthesis, which can be represented by the equation, 6 CO2 1 6 H2O : C6H12O6 1 6 O2.

14.45 3219 years

14.7 Rutherford examined the ability of radiation emitted from uranium to pass through thin aluminum foil. His key observation was that some particles emitted by the uranium went through foil but others were blocked by it. Rutherford concluded that there must be at least two types of particles. 14.9 Alpha decay: Atomic number decreases by 2, mass number decreases by 4. Positron decay: Atomic number decreases by 1. Gamma decay: No change in mass number or atomic number. K capture: Atomic number decreases by 1. Beta decay: Atomic number increases by 1. 14.11 14.13

(a) 209 F 184 4 (a) 188 83Bi : 81Tl 1 2He 40 40 K : 18 Ar 1 10b 1 n (c) 19

(b) 236 (c) 210 b 94 Pu _ 87 87 (b) 37Rb : 38Sr 1 210b 1 n 138 0 (d) 138 57Bi 1 21e : 56Ba 1 n

14.15 (a) Tritium has only three nucleons, so it cannot emit an alpha particle, which itself must contain four nucleons. (b) Beta emission (c) 14.17 (a) 197 79 Au

_ : 32 He 1 210 b 1 n 4 3 10 n : 197 79 Au 1 3 2 He

3 1H

(b) 206 82 Pb 1

1 3 210 b

2

14.19 4.0 3 10 years 14.21 (a) 0.0301 s21

(b) 1.7 3 1020 Bq, or 4.5 3 109 Ci

14.49 The heaviest element known to form a compound is fermium, which forms FmCl2. Presumably, heavier elements could form chemical compounds, but the small amounts formed in experiments, combined with the short half-lives of many isotopes, make observation of such compounds very difficult. 28 1 14.51 (a) 27 13 Al 1 1 p : 14 Si 1 g

_

41 0 1 (b) 40 18 Ar 1 0 n : 19 K 1 21 b 1 n

14.53 Because it does not have a charge, a neutron is not repelled by the charges on electrons or protons in the target atom. 14.55 7.6089 3 1013 J/kg of 235U

_

239 239 0 14.57 (a) 239 92 U* (b) 93 Np : 94 Pu 1 21 b 1 n

14.59 (a) 38 metric tons (b) The half-lives are 16.9 hours for 97 40 Zr and 2.49 s for 137 52 Te. So, virtually none of these isotopes would accumulate over the course of a year. Most of the daughter nuclei will decay either while still in the reactor core or while in short term storage at the power plant. 137

Te

14.25 9200 years

97

97

Nb

Mo

137

137

I

Xe

137

Cs

137

Ba

137

14.27 (a) Not useful; too little tritium left (b) Could be used; a modern wine is probably within a few half-lives (,60 or so years) old (c) Not useful; tomb objects are too old. 14.29 17,200 years 14.31 (a) 1 (b) 1.28 (c) 1.6 73 0 14.33 (a) Electron capture: 73 36 Kr 1 21 e : 35 Br 1 n

(b) Electron capture: 136 O 1 210 e : 135 N 1 n _

126 0 (c) Beta decay: 126 50 Sn : 51 Sb 1 21 b 1 n 232 4 228 90 Th : 88 Ra 1 2 He _ 228 228 0 88 Ra : 89 Ac 1 21 b 1 n _ 228 228 0 89 Ac : 90 Th 1 21 b 1n 224 4 228 90 Th : 88 Ra 1 2 He 224 4 220 88 Ra : 86 Rn 1 2 He 4 220 216 86 Rn : 84 Po 1 2 He 4 216 212 84 Po : 82 Pb 1 2 He _ 0 212 212 82 Pb : 83 Bi 1 21 b 1n _ 0 212 212 83 Bi : 84 Po 1 21 b 1 n 4 208 212 84 Po : 82 Pb 1 2 He

14.47 With an N/Z ratio of 1, 164Pb would lie far above the band of stability, making it susceptible to positron emission or electron capture.

(c) 97Zr

14.23 54 mg

14.35

14.43 1.67682 3 10211 J/nucleus, or 1.00981 3 1010 kJ/mol

The half-life of Cs is about 30 years. All the other isotopes involved have half-lives less than 1 day. (d) From the half-lives, we can conclude that virtually all of the 97Zr will have decayed to stable 97Mo before the waste would be disposed of. The 137Te will have decayed to 137Cs and will still be radioactive due to the longer half-life of that isotope. Thus we might estimate that the amount of radioactive waste actually would be 137/(97 1 137) of the original estimate, or 22 metric tons. This assumes that there is an easy and affordable way to separate the stable 97 Mo from the radioactive 137Cs. 14.61 (a) 115 B 1 11 p : 126 C 1 g (The gamma ray carries away excess energy.) (b) Boron-11 occurs naturally, whereas tritium would have to be produced artificially to carry out deuterium– tritium fusion. Also, the reaction considered here does not release neutrons, making it easier to contain radiation. 14.63 In the event of a nuclear plant emergency, there is a strong chance that radioactive iodine could be released. Because iodine accumulates in the thyroid, this could damage the thyroids of people who breathe the air nearby. Potassium

Appendix K

551

iodide tablets act by saturating the thyroid with normal iodine and preventing the gland from taking up appreciable amounts of radioactive iodine. 14.65 Designing a structure that must last for thousands of years poses many challenges. An engineer working on such a project must attempt to anticipate conditions far into the future. Dramatic climate change could occur on this time-scale, for example, and might impact structural requirements. Predicting human activity far in the future is even more difficult. Roads built in the days of horse-drawn carriages could not be expected to withstand modern traffic, for example. 14.67 The smallest ionization energies are on the order of 500 kJ/mol. This corresponds to a photon energy of about 10218 J. This means that ultraviolet light is likely to be the lowest energy ionizing radiation. 14.69 g > b > a 14.71 Inhaling the smoke brings 210Po into the lungs, and alpha particles emitted from inside the body are much more dangerous than those coming from external sources. 14.73 Modern electronic devices rely on the carefully controlled distribution and flow of electric charges. Cosmic rays can produce ions in these devices, disrupting their normal function. Chips for use in satellites are often encased in materials that are less susceptible to penetration by radiation to minimize damage by cosmic rays. 14.75 5.4 3 1027 Ci, or 0.54 mCi

552

Appendix K

14.77 (a) 7200 J

(b) 0.12 Gy

14.79 Because cosmic rays interact with gases in the atmosphere, natural radiation levels tend to increase with altitude. Radiation levels at the altitudes where commercial airliners fly are noticeably higher than those at ground level, and shielding planes to block all such radiation would be difficult and expensive. 14.81 (a) 189 F : 188 O 1 01 b 1 n 14.83

(c) 137 N : 136 C 1 01 b 0 0 1 b 1 21 b : 2 g

(b) 116 C : 115 B 1 01 b 1 n

1n

14.85 Responses will vary depending on the sources listed. The method is called dual photon absorptiometry, or DPA, and is based on the fact that when 153Gd decays it emits gamma ray photons at multiple energies. Comparing the intensities of the emission at two of these energies allows calculating bone density because calcium absorbs one of these photon energies more strongly than the other. 14.87 The most important piece of information needed is the half-life of 201Tl. It might also be useful to know how quickly the isotope was expected to be used once it had been delivered. 14.89 Relevant information that could be found easily includes the half-life and decay mode for each isotope. It would also be important to understand what type and energy of radiation was most likely to be effective in killing tumor cells. This information might be more difficult to look up.

Glossary A Absolute temperature (5.6): Temperature measured on a scale whose zero point is absolute zero. Kelvin is the most common absolute temperature scale, and Rankine is another example. Absorbed dose (14.7): The amount of any type of radiation absorbed in any kind of material. Acceptor level (8.3): Energy state associated with the dopant atoms and lying slightly above the valence band in a p-type semiconductor. Accuracy (1.3): How close the observed value is to the “true” value. Acid (3.3): Any substance that dissolves in water to produce H3O1 (or H1) ions. See also Brønsted–Lowry acid. Acid ionization constant (Ka) (12.7): Equilibrium constant for the reaction of a weak acid with water, producing H3O1 and the conjugate base of the weak acid. Actinides (2.5): The elements thorium through lawrencium, with valence electrons in the 5f subshell. They appear as the second row of elements below the main body of the periodic table. Activated complex (11.5): The arrangement of atoms corresponding to the maximum potential energy along the lowest energy path from reactants to products. Also called the transition state. Activation barrier (11.5): See activation energy. Activation energy (11.5): The minimum energy that reactant molecules must possess to undergo a particular chemical reaction. Active electrolysis (13.6): Electrolysis in which the electrode materials are reactants and products in the cell reaction. Actual yield (4.4): The amount of product obtained in a chemical reaction. Addition polymerization (8.6): Reaction in which monomers are added to one another to form a polymer, and no other products are formed. Addition reaction (2.6): Reaction in which two or more smaller molecules combine to form a single larger product molecule. Adhesion (8.5): Attractive forces between two dissimilar substances or surfaces. Admixture (12.1): Chemical substances added to concrete to give it certain desirable characteristics or properties. Alkali metals (2.5): Elements of Group 1, the first column of the periodic table, except hydrogen: Li, Na, K, Rb, Cs, and Fr. Alkaline battery (13.5): The most common single-use (primary) battery, so named because it uses a paste of KOH as the electrolyte. Alkaline earth metals (2.5): Elements of Group 2, the second column of the periodic table: Be, Mg, Ca, Sr, Ba, and Ra. Alkanes (4.1): Hydrocarbons in which the carbon atoms are all linked together by single bonds. Alpha decay (14.2): A radioactive decay process in which an atom’s nucleus emits an alpha particle ( 42 He) and that reduces the atomic number by two and the mass number by four. Alpha particle (14.2): An energetic particle with a charge of 21 and a mass number of 4 emitted in some radioactive decays. A helium nucleus.

Alpha ray (14.2): Radiation consisting of alpha particles. Alternating copolymer (8.6): Polymer in which molecules contain a regular alternating pattern of two different monomers. Amorphous (8.2): Noncrystalline solid in which the arrangement of atoms or molecules has no long-range order. Amplitude (6.2): The size or “height” of a wave. Anion (2.3): A negatively charged atom or group of atoms. Anode (13.2): Electrode at which oxidation takes place. Antibonding molecular orbital (8.3): An orbital resulting from the combination of atomic orbitals from two or more atoms that has higher energy than any of the atomic orbitals from which it was formed. Aqueous solution (3.3): Solution in which water is the solvent. Arrhenius behavior (11.5): Variation of a rate constant with temperature in accordance with the Arrhenius equation. Arrhenius equation (11.5): Mathematical description of the dependence of the rate constant on temperature: k 5 Ae Ea/RT. Atactic polymer (8.6): A polymer in which substituent groups attached to the polymer chain are not arranged in a regular or ordered fashion. Atmosphere (atm) (5.2): Unit of pressure equal to 760 torr or 101,325 Pa. Normal atmospheric pressure at sea level is approximately 1 atm. Atomic mass unit (amu or u) (2.2): Unit used to express the relative masses of atoms. The amu is defined as one-twelfth the mass of an atom of 12C, and 1 amu 5 1.6606 3 10224 g. Atomic number (Z ) (2.2): The number of protons in the nucleus of an atom of a particular element. Atomic spectrum (6.3): The pattern of wavelengths of light absorbed or emitted by a particular element. Atom (1.2): An unimaginably small particle that could not be made any smaller and still behave as a chemical system. Atoms are the smallest particles that can retain an element’s chemical identity. Aufbau principle (6.5): Description of the process by which electrons fill orbitals in order of increasing energy. Average rate (11.2): Rate of reaction measured over some period of time. Average speed (5.6): The mean speed of molecules in a gas; can be found from the Maxwell-Boltzmann distribution of speeds. Avogadro’s law (5.3): At fixed temperature and pressure, the volume of a gas is proportional to the number of molecules (or moles) of gas present. Avogadro’s number (NA) (3.4): The number of particles in a mole of any substance (6.022 3 1023).

B Background radiation (14.7): Radiation originating from both natural and man-made sources to which an entire population is routinely exposed. Band gap (8.3): Difference in energy between the valence band and the conduction band in a solid.

553

Band of stability (4.4): A region in the chart of the nuclides in which all known stable nuclides are located. The band of stability lies near N/Z 5 1 for light nuclei but curves toward higher N/Z ratios for heavier elements. Band theory (8.3): Description of chemical bonding in solids in terms of bands of orbitals. Used to explain the electrical properties of metals, insulators, and semiconductors. Barometer (5.2): A device that measures atmospheric pressure. Base (3.3): Any substance that dissolves in water to produce OH2 ions. Base ionization constant (Kb) (12.7): Equilibrium constant for the reaction of a weak base with water, producing OH2 and the conjugate acid of the weak base. Battery (13.5): An electrochemical cell or series of cells in which a spontaneous redox reaction generates an electric current. Becquerel (Bq) (14.3): The SI unit of nuclear activity, equal to one nuclear disintegration per second. Beta decay (14.2): A radioactive decay process in which a neutron decays into a proton and an electron (or beta particle, 210 b). The beta particle and an antineutrino are emitted while the proton remains in the nucleus, thus increasing the atomic number by one and leaving the mass number unchanged. Beta particle (14.2): An energetic particle with a charge of 1– and a mass number of 0 emitted in some radioactive decays. An electron emitted from the nucleus. Beta ray (14.2): Radiation consisting of beta particles. Bimolecular (11.6): Describes an elementary step involving a collision between two reactant molecules. Binary compound (2.7): Any chemical substance containing two different elements. Binding energy (6.2): Amount of energy required to remove an electron from a metal surface. Also called the work function. See also nuclear binding energy. Biocompatibility (7.1): Ability of a material to coexist with a natural biological system into which it has been implanted. Block copolymer (8.6): Polymer in which molecules contain long sections or blocks of two or more different monomers. Body-centered cubic (bcc) (8.2): Crystal lattice in which particles are arranged at the corners of a cube and also at the center of the cube. Bohr model (6.3): Early model of atomic structure that proposed that electrons are restricted to circular orbits of specific sizes. Although this model did account for some features of atomic spectra, it is no longer viewed as correct. Bond energy (7.3): The amount of energy released when isolated atoms form a covalent bond. Also equal to the amount of energy that must be supplied to break that bond. Bond length (7.3): The distance between the nuclei of two bonded atoms. Bonding molecular orbital (8.3): An orbital resulting from the combination of atomic orbitals from two or more atoms that has lower energy than any of the atomic orbitals from which it was formed. Bonding pair (7.3): A pair of electrons shared in a covalent bond between two atoms. Boundary (9.3): The line or surface—often hypothetical—that separates the system and the surroundings in a thermodynamic problem.

554

Glossary

Boyle’s law (5.3): At constant temperature, the pressure and volume of a fixed amount of gas are inversely proportional. British thermal unit (Btu) (9.1): A unit of energy defined as the amount of energy needed to raise the temperature of 1 lb of water by 1°F. Brønsted–Lowry acid (12.7): Any species that donates a proton (H1 ion) in a chemical reaction.

C Calorie (cal) (9.2): A unit of energy originally defined as the amount of energy required to heat 1 g of water from 14.5 to 15.5°C. 1 cal 5 4.184 J. (Note that a “food calorie,” abbreviated “Cal,” is actually a kilocalorie. 1 Cal 5 4.184 kJ) Calorimetry (9.4): Experimental technique for measuring the heat flow into or out of a thermodynamic system. Catalyst (11.7): Any substance that increases the rate of a chemical reaction but is neither created nor destroyed in the process. Cathode (13.2): Electrode at which reduction takes place. Cathodic protection (13.8): A method for preventing corrosion whereby the metal to be protected is connected electrically to a more reactive metal. The protected metal thus becomes the cathode, so it cannot be oxidized. Cation (2.3): A positively charged atom or group of atoms. Cell notation (13.2): Shorthand notation for an electrochemical cell that shows the electrodes, gases, and solutions in the cell reaction as well as the phase boundaries separating them. Cell potential (13.2): The electromotive force of an electrochemical cell. The amount of work the cell can do per coulomb of charge. Chain reaction (14.6): A self-sustaining series of reactions. In nuclear fission, a situation in which one or more neutrons released in one fission event induce further fission. Chapman cycle (11.1): A series of chemical reactions responsible for the formation and destruction of ozone in the stratosphere. Charles’s law (5.3): The volume and temperature of a fixed amount of gas at constant pressure are directly proportional. Chart of the nuclides (14.4): A plot of atomic number (Z) versus neutron number (N) for all known nuclei. Chemical bond (2.4): Attractive force between two atoms holding them together to form a molecule or a part of a molecule. Chemical compound (2.4): A pure substance made up of atoms of two or more elements joined together by chemical bonds. Chemical energy (9.2): Potential energy stored in chemical bonds, which can be released through an exothermic reaction. Chemical equation (3.2): Symbolic representation of a chemical reaction. Chemical formula (2.4): Symbolic description of a molecule or compound showing its constituent elements and the relative number of each. Chemical nomenclature (2.7): The systematic assignment of names to chemical compounds on the basis of their constituent elements or ions and their structures. Chemical property (1.2): Property of a substance that is associated with the types of chemical changes that the substance undergoes. Chemical reaction (3.2): Process in which substances (reactants) are transformed into other substances (products) by the combination, rearrangement, or separation of atoms. Coherent light (6.8): Light in which all waves are in phase, such as the light produced by a laser.

Cohesion (8.5): Attractive forces between two similar substances or surfaces. Combustion (1.2): Reaction in which an element or compound burns in oxygen or air. Complete combustion of hydrocarbons produces CO2 and H2O as products. Common ion effect (12.6): Shift in equilibrium when one or more ions that are part of the equilibrium are introduced from an outside source. Compound nucleus (14.6): A high-energy intermediate formed in a nuclear reaction, similar to an activated complex in a chemical reaction. Concentration (3.3): A measure of the relative quantities of solute and solvent in a solution. Condensation (9.5): The process by which a molecule from the gas phase enters the liquid phase. Also the process by which a bulk gas is converted to a liquid. Condensation polymer (8.6): A polymer made from monomers containing two or more functional groups through a reaction in which water (or another small molecule) is formed as a byproduct. Conduction band (8.3): A band of orbitals or energy states at higher energy than the valence band and typically unoccupied in the ground state of a solid. Conjugate acid (12.7): Molecule or ion formed by the addition of a hydrogen ion to its conjugate base. Conjugate acid–base pair (12.7): Two chemical species related to one another by the loss or gain of a single hydrogen ion. Conjugate base (12.7): Molecule or ion formed by the removal of a hydrogen ion from its conjugate acid. Control rod (14.6): A rod inserted into a nuclear reactor to regulate the rate of fission. Control rods are made from elements that can absorb neutrons without undergoing subsequent decay. Coordination number (8.2): The number of lattice points immediately adjacent to any given lattice point in the structure of a solid. Copolymer (8.6): A polymer made from more than one type of monomer. Core electrons (6.5): Electrons in the filled inner shell(s) of an atom. Corrosion (13.1): The degradation of metals exposed to the environment by chemical oxidation. Cosmic ray (14.1): Energetic particles impinging upon the atmosphere from beyond Earth. Coulomb’s law (2.3): Mathematical description of the force between two charged particles. Covalent bond (2.4): Attractive force resulting from the sharing of electrons between pairs of atoms. Crevice corrosion (13.1): Corrosion in a small gap between two pieces of metal. Criteria pollutants (5.1): Chemical compounds found in urban air that have a variety of negative effects on health, the environment, and property. The six principal criteria pollutants identified by the EPA are CO, NO2, O3, SO2, Pb, and particulate matter (PM). Critical mass (14.6): The minimum mass of a particular fissile fuel that can sustain a chain reaction. Crystalline solid (8.2): Solid featuring regular, repeating geometric arrangements of atoms, ions, or molecules. Curie (Ci) (14.3): A unit of nuclear activity, exactly equal to 3.7 3 1010 Bq.

D Dalton’s law of partial pressures (5.4): The total pressure of a mixture of gases is the sum of the partial pressures of the component gases. Decay constant (14.3): The first-order rate constant for a radioactive decay. We use the symbol k in this text, but l is also commonly used. Decay series (14.4): A series of nuclear decays starting with an unstable isotope and eventually producing a stable nuclide. Deductive reasoning (1.3): A thought process that takes two or more statements or assertions and combines them so that a clear and irrefutable conclusion can be drawn. Degree of polymerization (8.6): The average number of repeating units in a polymer. Density ( r) (1.2): The ratio of mass to volume. Differential rate law (11.3): Mathematical expression for the rate of a chemical reaction defined in terms of the rate of change in the concentration of a reactant or product. Diffraction (6.4): Interaction of a wave with a regular pattern of lines or slits. Dilution (3.5): Process in which solvent is added to a solution to decrease the solute concentration. Dimensional analysis (1.5): Problem-solving strategy in which one inspects the units on all quantities in a calculation to check for correctness. Dipole (7.4): Two equal and opposite electric charges separated by a short distance. Dipole–dipole forces (8.4): Attractive interaction between two polar molecules or between two polar regions of a single large molecule. Dispersion forces (8.4): Attractive forces between molecules resulting from the interaction of transient (induced) dipoles. Also called London forces or van der Waals forces. Dissociation reaction (3.3): Chemical reaction in which an ionic solid dissolves in water and dissociates into its constituent ions. Distribution function (5.6): Mathematical function describing the variation of some property across a population. Donor level (8.3): Energy state associated with dopant atoms and lying slightly below the conduction band in an n-type semiconductor. Doping (8.3): The controlled addition of a very small concentration of one element (the dopant) to an intrinsic semiconductor such as silicon or germanium. Dosimeter (14.7): Any device used to measure an individual’s cumulative exposure to radiation. Double bond (7.3): A covalent bond formed by the sharing of two pairs of electrons between the same pair of atoms. Ductile (8.3): Easily drawn or pulled into a wire. Dynamic equilibrium (8.5, 12.2): State achieved when the rates of the forward and backward processes in a chemical reaction become equal to one another.

E Effective nuclear charge (Zeff) (6.5): Net positive charge “experienced” by outer shell electrons in a many-electron atom. Elastic modulus (1.6): A measure of the stiffness of a material, defined as the ratio of stress to strain. Electrochemistry (13.2): The study of the connection between oxidation–reduction reactions and the flow of electrons. Electrode (13.2): An electrically conducting site at which either oxidation or reduction takes place.

Glossary

555

Electrolysis (13.6): Process in which an electric current is used to drive a redox reaction in its nonspontaneous direction. Electrolyte (3.3): A substance that ionizes or dissociates in water to produce an aqueous solution that conducts electricity. Electromagnetic spectrum (6.2): The various forms of light, all of which consist of oscillating electric and magnetic fields traveling at the speed of light (2.998 3 108 m/s). Electromotive force (EMF) (13.2): The difference in electric potential between the two electrodes in a galvanic cell. See also cell potential. Electron affinity (6.7): The amount of energy needed to add an electron to a neutral atom to form an anion. Electron capture (14.2): A radioactive decay process in which an atom’s nucleus captures an inner shell electron, reducing the atomic number by one and leaving the mass number unchanged. Also called K capture. Electron configuration (6.5): A complete description of the distribution of electrons among orbitals in an atom or ion. Electronegativity (7.4): A relative measure of the tendency of an atom in a molecule to attract the shared electrons in a covalent bond. Electrons (2.2): Negatively charged subatomic particles that occupy most of the volume of an atom. Electron volt (eV) (14.1): A unit of energy equal to 96.5853 kJ mol–1. Electroplating (13.6): The process of depositing a thin coat of a metal on a surface by electrolysis. Elemental analysis (3.5): Experimental measurement of the mass percentage of each element in a compound. Elementary step (11.6): Chemical reaction taking place through a single collision between molecules, ions, or atoms. A series of elementary steps makes up a reaction mechanism. Element (1.2): A substance that cannot be broken down into two or more new substances by chemical or physical processes. Empirical formula (2.4): Chemical formula that shows the composition of a compound in terms of the simplest wholenumber ratio of the atoms of each element. Endothermic (9.5): A process in which heat flows into a thermodynamic system. For a chemical reaction, this means DH is positive. Energy density (9.7): The amount of energy that can be released per gram of fuel burned. Energy economy (9.1): The production and use of energy in the developed world. Enthalpy (DH ) (9.5): A state function defined as E 1 PV and equal to the heat flow at constant pressure. Enthalpy diagram (9.6): Conceptual device for visualizing the enthalpy changes in a chemical reaction. Enthalpy driven (10.6): Description of a process for which DH and DS are both negative and which is spontaneous only at low temperatures. Entropy (DS ) (10.3): A state function that measures the dispersal of energy in a system. Entropy is correlated with the degree of randomness or disorder of a system on the molecular level. Entropy driven (10.6): Description of a process for which DH and DS are both positive and which is spontaneous only at high temperatures. Equilibrium constant (Keq, K) (12.3): The ratio of concentrations of products to those of reactants; each concentration is raised to a power that matches its stoichiometric coefficient

556

Glossary

in the balanced chemical equation. Equilibrium constants vary with temperature. Equilibrium expression (12.3): Equation defining the equilibrium constant in terms of the concentrations of the various substances involved in the reaction. Equivalent dose (14.7): A measure of radiation exposure that attempts to measure the tissue damage likely to be done by a particular type of radiation. Equal to the absorbed dose multiplied by a quality factor. Excited state (6.3): Any state formed when an atom or molecule absorbs energy, as when an electron is promoted to a higher energy orbital. Exothermic (9.5): A process in which heat flows out of a thermodynamic system. For a chemical reaction, this means that DH is negative. Exposure (14.7): A measure of radiation exposure based on the ionization of molecules of air, expressed in coulombs per kilogram of air.

F Face-centered cubic (fcc) (8.2): Crystal lattice in which particles are arranged at the corners of a cube and also at the center of each face of the cube. Also called cubic close packing. Factor-label method (1.5): See dimensional analysis. Faraday constant (F ) (13.3): The electric charge on one mole of electrons: 96,485 C mol21. First law of thermodynamics (9.3): Energy can be transformed from one form to another but cannot be created or destroyed; the total energy of the universe is constant. First-order reaction (11.3): Any reaction whose rate is given by a rate law of the form, Rate 5 k[A], where A represents a chemical species in the reaction (usually a reactant). Fissile (14.6): Able to undergo nuclear fission. Also called fissionable. Formation reaction (9.5): A chemical reaction by which one mole of a compound is formed from its constituent elements in their standard states. Formula unit (2.4): The smallest whole number ratio of anions and cations in an ionic compound. Free radicals (2.8): Extremely reactive chemical species containing one or more unpaired electrons. Freezing (9.5): Phase change in which a liquid is converted to a solid. Frequency (n) (6.2): The number of complete cycles of a wave passing a given point per unit time. Frequency factor (11.5): The preexponential term (A) in the Arrhenius equation. This term accounts for the frequency of collisions as well as any geometric or orientational constraints that influence the reaction rate. Also called the preexponential factor. Fuel additives (4.6): Substances added to gasoline or other fuels, typically intended to improve performance or reduce exhaust emissions. Fuel cell (13.5): An electrochemical cell in which reaction between a fuel and an oxidant produces electricity. Unlike a battery, a fuel cell uses a replenishable external supply of reactants. Fuel rod (14.6): Component of a nuclear reactor containing fissionable uranium fuel. Functional group (2.6): An atom or group of atoms that imparts characteristic chemical properties to any organic compound in which it is found. Fusion (9.5): The phase change in which a solid is converted to a liquid; melting. See also nuclear fusion.

G Galvanic cell (13.2): Any electrochemical cell in which a spontaneous chemical reaction generates an electric current. Galvanic corrosion (13.1): Degradation of metal parts through oxidation that occurs only when two different metals are in contact with each other in the presence of an appropriate electrolyte. Galvanized steel (13.3): Steel coated with a thin layer of zinc to prevent corrosion. Gamma ray (14.2): Radiation consisting of high-energy photons. Gas (1.2): State of matter in which a substance has no fixed shape and expands to occupy the entire volume of its container. Geiger counter (14.7): A radiation detector generally sensitive only to alpha and beta radiation. Also called a Geiger-Mueller counter. Gibbs free energy (G ) (10.6): A thermodynamic state function that can be used to predict the direction of spontaneous change at constant temperature and pressure. DG 5 DH 2 TDS, and DG , 0 for any spontaneous process. Graft copolymer (8.6): Material whose molecules contain large segments of one type of molecule “grafted” onto or attached to chains of a second type of polymer. The graft copolymer will display properties of each of the constituent polymers. Green chemistry (3.6): An approach to chemical synthesis and processes to minimize the production of waste and pollutants. Ground state (6.3): The lowest energy state of an atom or molecule. Group (2.5): Vertical column of the periodic table.

H Half-life (t1/2) (11.4): The time it takes for the concentration of a reactant to fall to one-half of its original value. For a first-order reaction, the half-life is independent of the initial concentration. Half-reaction (13.2): A reaction showing only the oxidation or reduction process in a redox reaction. Halogens (2.5): The elements in Group 17 of the periodic table: F, Cl, Br, I, and At. HDPE (2.8): High-density polyethylene. Heat (q) (9.2): The flow of thermal energy between two objects, from the warmer to the cooler because of the difference in their temperatures. Heat capacity (9.4): The amount of energy that must be transferred to an object or substance to raise its temperature by a specified amount (usually 1 K). Heat of formation (DHf°) (9.5): Enthalpy change for the formation of one mole of a substance from its constituent elements in their thermodynamic standard states. Heat of reaction (DH °) (9.5): Enthalpy change for a chemical reaction. Heisenberg’s uncertainty principle (6.4): It is impossible to determine both the position and momentum of an electron simultaneously and with arbitrary accuracy. Hess’s law (9.6): The enthalpy change for any process is independent of the way the process is carried out. Heterogeneous catalyst (11.7): Catalyst in a phase different from the reacting substances. Heterogeneous equilibrium (12.3): Chemical equilibrium in which not all reactants and products are in the same phase.

High-density polyethylene (10.8): Form of polyethylene featuring long, straight chain molecules. Because these chains pack together more tightly, HDPE is harder and stronger than low-density polyethylene. Homogeneous catalyst (11.7): Catalyst in the same phase as the reacting substances. Homogeneous equilibrium (12.3): Chemical equilibrium in which all reactants and products are in the same phase. Hund’s rule (6.5): Within a subshell of an atom, electrons occupy orbitals individually first before pairing. Hybrid orbitals (7.7): Orbitals formed by combining two or more orbitals from the same atom. Hybridization (7.7): The process of forming hybrid orbitals. The idea that when two or more atoms approach one another closely enough to form chemical bonds, the interactions can be strong enough to reshape the orbitals of those atoms. Hydrates (2.4): Compounds in which one or more water molecules are incorporated into the crystal lattice. Hydrocarbons (2.6): Molecules containing only carbon and hydrogen atoms. Hydrogen bonding (8.4): An unusually strong dipole–dipole interaction between a hydrogen atom bonded to F, N, or O and a lone pair of electrons on another highly electronegative atom such as F, N, or O.

I Indicator (4.5): A dye added to a titration to show when a reaction is complete. Inductive reasoning (1.3): Thought process that begins with a series of specific observations and attempts to generalize to a larger, more universal conclusion. Inorganic chemistry (2.6): The study of all elements, besides carbon, and their compounds. Insoluble (3.3): Description of a solid that does not dissolve in solution to any significant extent. Instantaneous rate (11.2): The rate of a chemical reaction measured at a single instant. Insulator (8.3): A material that does not conduct electricity, characterized by a large energy gap between its valence and conduction bands. Integrated rate law (11.4): Mathematical equation that shows how the concentration of a reactant or product changes as a function of time. Intermolecular forces (8.4): Attractive forces between molecules. Internal energy (E ) (9.2): The combined kinetic and potential energies of the atoms and molecules that make up a substance or object. Ion (2.3): An atom or group of atoms with a net electrical charge. Ionic bonds (7.2): Chemical bonds formed by the mutual electrostatic attraction of oppositely charged ions. Ionic bonding (2.4): Bonding in which one or more electrons are transferred from one species to another. Ionization energy (6.7): The amount of energy required to remove an electron from a free atom in the gas phase. Ionization gauge (5.7): A pressure sensor in which gas molecules are first ionized and then collected and measured as a current. Suitable for measurement of low pressures, between 1025 and 10211 torr.

Glossary

557

Ionizing power (14.7): A measure of the ability of a particular type of radiation to produce ions in matter with which it interacts. Irreversible (10.6): A process, often far from equilibrium, which cannot be reversed by a small incremental change in any variable. Isomers (4.1): Compounds having the same molecular formula but in which the atoms are arranged differently. Isotactic polymer (8.6): A polymer in which substituent groups are all on the same side of the polymer chain. Isotopes (2.2): Atoms of the same element that have different numbers of neutrons. Isotopic abundance (2.2): The percentage of the atoms in a naturally occurring sample of an element that consists of a specified isotope.

J Joule ( J) (9.2): The SI unit of energy; 1 J 5 1 kg m2/s2

K K capture (14.2): A radioactive decay process in which an atom’s nucleus captures an inner-shell electron, reducing the atomic number by one and leaving the mass number unchanged. Also called electron capture. Kinetic energy (9.2): Energy from the motion of an object; given by —21 mv2, where m is the mass of the object and v is its velocity. Kinetic–molecular theory (5.6): Description of a gas as a collection of a very large number of atoms or molecules in constant, random motion. The ideal gas law can be derived from the postulates of the kinetic theory.

L Lanthanides (2.5): The elements cerium through lutetium, whose valence electrons are in the 4f subshell. They appear as the first row of elements below the main body of the periodic table. Laser (6.8): A source of high-intensity, coherent, monochromatic light. Originally an acronym for “l ight amplification by stimulated emission of radiation.” Lattice (2.4, 7.2): A regular, periodic arrangement of atoms, ions, or molecules in a crystalline solid. Lattice energy (7.2): Enthalpy change from the formation of a crystal lattice from its separate constituent ions. Law (1.3): A statement of an accepted scientific principle, based on evidence from a significant number of experimental observations. Law of conservation of matter (3.2): Matter is neither created nor destroyed in ordinary chemical reactions. LDPE (2.8): Low-density polyethylene Lead-acid storage battery (13.5): Secondary (rechargeable) battery used in the ignition system of automobiles, whose name derives from the use of lead in the electrodes and sulfuric acid as the electrolyte. LeChatelier’s principle (12.5): When a system at equilibrium is stressed, it responds by reestablishing equilibrium so that the applied stress is partially offset. Lewis dot symbol (7.3): Visual representation of the valence electrons of an atom. Lewis structure (7.3): Visual representation of the structure of a covalently bonded molecule or ion that shows the distribution of all valence electrons. Light-emitting diode (LED) (6.8): Semiconductor device in which monochromatic light is produced as electrons undergo a transition across the band gap.

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Glossary

Limiting reactant (4.3): The reactant that is completely consumed in a reaction. The available amount of it determines the maximum possible reaction yield. Line structure (2.6): Visual representation of organic compounds in which carbon atoms and the hydrogen atoms attached to them are not shown explicitly. Liquid (1.2): State of matter in which a substance has a fixed volume but flows to take on the shape of its container. Lone pair (7.3): Paired valence electrons that are associated with a single atom and not involved in a covalent bond; also called nonbonding electrons. Low-density polyethylene (10.8): Form of polyethylene featuring branched chain molecules. Because these chains do not pack together as tightly, LDPE is softer and more flexible than high-density polyethylene.

M Macroscopic perspective (1.2): Viewpoint of chemistry focusing on samples of matter that are large enough to be seen, measured, or handled easily. Magic number (14.5): A number of neutrons or protons associated with unusually stable nuclei. Known magic numbers are 2, 8, 20, 28, 50, 82, 126, and 184. Magnetic quantum number (m,) (6.4): One of the three quantum numbers identifying an atomic orbital. Specifies the spatial orientation of an orbital. Allowed values of m, are related to the value of the secondary quantum number (,). Main group elements (2.5): Elements of Groups 1, 2, and 13218, for which the highest lying valence electron occupies an s or p orbital. Also called representative elements. Malleability (1.2): Ability of a material to be shaped or formed by the application of pressure, as through hammering or rolling. Malleable (8.3): Able to be shaped or formed by the application of pressure, by hammering or rolling, for example. Manometer (5.7): A device used to measure pressure. Mass defect (14.5): The difference in mass between a nucleus and the sum of the masses of the neutrons and protons that make up the nucleus. Related to nuclear binding energy by Einstein’s equation. Mass density (1.3): The ratio of the mass of a substance to its volume. Also known simply as density. Mass number (2.2): The combined total of protons and neutrons in an atom. Matter (1.2): Anything that has mass and occupies space. Maxwell-Boltzmann distribution (5.6): Mathematical function describing the fraction of molecules traveling at a particular speed in a gas. Depends on the mass of the molecules and the temperature of the gas. Mean free path (5.6): The average distance a particle travels between collisions with other particles. Mercury battery (13.5): Commercial battery in which zinc is oxidized and mercury(II) oxide is reduced. Used in calculators, watches, and hearing aids. Metal (2.5): An element that is malleable, ductile, conducts electricity, and tends to form cations in its chemical compounds. Metallic bonding (2.4): Bonding in a solid metal, where the nuclei and some fraction of their electrons comprise a positively charged “core” localized at these lattice points, and valence electrons form a surrounding “sea” of negative charge. Metalloid (2.5): An element that displays some properties of metals and some properties of nonmetals. Sometimes called a semimetal.

Methyl tertiary-butyl ether (MTBE) (4.6): Compound with the formula C5H12O that has been used as an oxygenating additive in gasoline. Microscopic perspective (1.2): Viewpoint of chemistry focusing on samples of matter at the atomic and molecular level, where samples cannot be seen, measured, or handled easily. Note that this scale is smaller than the resolution of a traditional microscope. Also called the particulate perspective. Microstate (10.3): One possible way for a collection of particles to assume a given total energy. Molar concentration (3.5): See Molarity. Molar heat capacity (C ) (9.4): The amount of heat required to raise the temperature of one mole of a substance by one degree (°C or K). Molar mass (3.4): The average mass of one mole of any element or compound. Molar solubility (12.6): The concentration of a solid in a saturated solution, expressed in moles per liter. Molarity (M) (3.5): Unit of concentration defined as the number of moles of solute per liter of solution. Mole (mol) (3.4): The particular quantity by which chemists usually choose to count atoms or molecules. The number of particles in one mole is defined as the number of atoms in exactly 12 grams of 12C, which is also called Avogadro’s number. Mole fraction (X ) (5.4): The ratio of the number of moles of one component to the total number of moles in a mixture. Mole ratio (3.4, 4.2): A stoichiometric ratio relating the number of moles of one reactant or product in a chemical reaction to the number of moles of some other reactant or product in the same reaction. Obtained from a balanced chemical equation. Molecular equation (3.3): Chemical equation in which all substances are shown as molecular formulas. Molecular formula (2.4): Chemical formula showing the number of atoms of each element in one molecule of a compound. Molecular shape (7.8): The geometry of a molecule or ion, defined by the positions of its constituent atoms. Molecule (1.2, 2.4): Two or more atoms joined together by chemical bonds, forming the smallest particle that retains the chemical properties of a substance. Monochromatic (6.8): Light consisting only of a single wavelength or color, such as the light emitted by a laser. Monomer (2.1): The small repeating unit from which a polymer molecule is made. Most probable speed (5.6): The speed at which the most molecules move in a gas; the maximum in the MaxwellBoltzmann distribution of speeds.

N n-type semiconductor (8.3): Material formed by doping silicon or germanium with an element such as phosphorus or arsenic that adds extra valence electrons. So named because the charge carriers in an n-type semiconductor are negatively charged electrons. Nanotube (8.1): A long, cylindrical carbon molecule. Nanotubes hold promise for a wide range of future engineering applications. Nernst equation (13.3): Mathematical equation relating the cell potential to the concentrations and pressures of chemical species in the cell reaction. Net ionic equation (3.3): A chemical equation in which those ions and molecules actually involved in the reaction are shown, but any spectator ions are omitted.

Neutralization (3.3): A chemical reaction in which an acid and a base react to produce water and a salt. Neutrons (2.2): Electrically neutral subatomic particles found in an atom’s nucleus. Nickel-cadmium (nicad) battery (13.5): Secondary (rechargeable) battery used in many small appliances; involves the oxidation of cadmium and the reduction of nickel oxyhydroxide (NiO(OH)). Nickel-metal-hydride (NiMH) battery (13.5): Secondary (rechargeable) battery used in many small electronic devices; involves the oxidation of hydrogen from a metal hydride and the reduction of nickel oxyhydroxide (NiO(OH)). Compared with nicad batteries, NiMH eliminates the use of toxic cadmium and is less prone to memory effects. Noble gases (2.5): The elements in Group 18 of the periodic table: He, Ne, Ar, Kr, Xe, and Rn. Node (6.4): A point or plane where the amplitude of a wave is equal to zero. For atomic orbitals, a point or plane where there is zero probability of finding an electron. Nonattainment area (5.1): A region where detected concentrations of one or more criteria pollutants exceed the limits specified in the EPA’s primary standards. Nonelectrolyte (3.3): A substance that dissolves in water to produce a solution that does not conduct electricity. Nonmetal (2.5): An element that does not display the chemical and physical properties of metals. Nuclear binding energy (14.5): The amount of energy released by the formation of an atomic nucleus from isolated nucleons. Also the amount of energy required to separate a nucleus into its constituent nucleons. Normal boiling point (8.5): The temperature at which a substance’s vapor pressure is equal to 1 atmosphere. Nuclear fission (14.5): A process in which a single heavy nucleus splits into two or more smaller and more stable nuclei. Nuclear fusion (14.5): A process in which two or more smaller nuclei combine to form a single heavier and more stable nucleus. Nuclear reaction (14.1): Any process in which an atom or nucleus is converted into a different chemical element. Nucleus (2.2): The small, dense core of an atom, containing neutrons and protons. Nuclide (14.2): A specific nucleus with the numbers of protons and neutrons specified. The nucleus of a particular isotope.

O Octet rule (7.3): In forming chemical bonds, main group elements gain, lose, or share electrons to achieve a configuration in which they are surrounded by eight valence electrons. Operator (6.4): Notation used to indicate one or more operations to be carried out on a mathematical function. Orbital (6.4): The wave representation of electrons in an atom; a mathematical or visual depiction of a region of space in which there is a high probability of finding an electron. Orbital overlap (7.6): Interaction between orbitals—and hence electrons—from two different atoms responsible for the formation of a covalent bond. Order of reaction (11.3): The way in which a reaction rate depends on the concentration(s) of the substance(s) involved, expressed as the exponents on the concentration term(s) in the rate law. Typically a small integer or a small rational fraction. Organic chemistry (2.6): The study of compounds of the element carbon.

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559

Oxidation (13.2): The loss of electrons by some chemical species in a chemical reaction. Oxidation–reduction reaction (9.8, 13.2): Chemical reaction involving the transfer of electrons from one species to another. Also called a redox reaction. Oxyanions (2.7): Polyatomic anions that contain oxygen. Oxygenated fuels (4.6): Blends of gasoline or other fuels to which oxygen-containing additives such as methanol, ethanol, or MTBE have been added to reduce undesirable exhaust emissions.

P p-n junction (8.3): Point at which a piece of p-type material meets a piece of n-type material; an important building block for integrated circuits. p-type semiconductor (8.3): Material formed by doping silicon or germanium with an element such as aluminum or gallium that reduces the number of valence electrons. So named because the charge carriers in a p-type semiconductor can be thought of as positively charged vacancies or “holes.” Packing efficiency (8.2): The percentage of the total volume that is actually occupied by the atoms in a given crystal lattice. Partial pressure (5.4): The pressure that would be exerted by one component of a gas mixture if only that component were present at the same temperature and volume as the mixture. Particulate perspective (1.2): Viewpoint of chemistry focusing on samples of matter at the atomic and molecular level, where samples cannot be seen, measured, or handled easily. Also called the microscopic perspective. Parts per billion (ppb) (1.4): A unit of concentration defined as the number of particles of a particular component per every billion (109) molecules of a mixture. Often used to describe the levels of pollutants or other trace components in air. Parts per million (ppm) (1.4, 5.1): A unit of concentration defined as the number of particles of a particular component per every million (106) molecules of a mixture. Often used to describe the levels of pollutants or other trace components in air. Pascal (Pa) (5.2): The SI unit for pressure. 1 Pa 5 1 N m22 Passivation (13.8): A process in which oxidation of a metal surface produces a protective layer that is not susceptible to further corrosion. Passive electrolysis (13.6): Electrolysis in which the electrodes are chemically inert materials that simply provide a pathway for electrons to enter and leave the electrolytic cell. Pauli exclusion principle (6.5): No two electrons in an atom may have the same set of four quantum numbers. Penetrating power (14.7): A measure of the depth to which a particular type of radiation is expected to penetrate in matter with which it interacts. Percentage yield (4.4): The ratio of the actual yield to the theoretical yield in a particular chemical reaction, expressed as a percentage. Periodic law (2.5): When the elements are arranged by atomic number, they display regular and periodic variation in their chemical properties. Periodicity (2.5): Regular patterns in the chemical and physical behavior of groups of elements. Periods (2.5): Horizontal rows of the periodic table. Phase diagram (8.1): A diagram showing what state of an element is most stable under different combinations of temperature and pressure.

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Phases (1.2): One of the three states of matter: solid, liquid, or gas. Photochemical reaction (3.2, 5.1): A chemical reaction in which light provides the energy needed to initiate the reaction. Photochemical smog (11.8): Common urban air condition characterized by a brownish haze; produced by a complex series of reactions involving ozone (O3), oxides of nitrogen (NOx), and volatile organic compounds (VOCs). Photoelectric effect (6.2): Emission of electrons from metals illuminated by light of an appropriate wavelength. Photon (6.2): A massless “particle” of light, whose energy is equal to Planck’s constant times the frequency of the light (hn). Physical property (1.2): Property that can be observed or measured while the substance being observed retains its composition and identity. Pi (p) bond (7.6): A chemical bond in which the electron density is not concentrated along the line of centers of the two atoms, as occurs through the “sideways” interaction of two p orbitals. Polar bond (7.4): A chemical bond between two atoms that have different electronegativities, such that one end of the bond takes on a partial positive charge and the other end takes on a partial negative charge. Also called a polar covalent bond. Polarizability (8.4): Susceptibility of a molecule to perturbations of its charge distribution in the presence of an external electric field. Poly(ethylene terephthalate) (PET) (10.8): A recyclable polyester commonly used in bottles for soft drinks and other liquids. Poly(vinyl chloride) (PVC) (10.8): A thermoplastic polymer made from vinyl chloride monomer; widely used in construction materials, pipes, and for wire and cable insulation. Polyatomic ion (2.7): A group of atoms that carries an electrical charge. Polymer backbone (2.1): The long chain of atoms—usually carbon atoms—running the length of a polymer molecule. Polypropylene (10.8): A thermoplastic polymer made from propylene monomer; widely used in syringes and other medical supplies because it can be sterilized easily. Polystyrene (10.8): A thermoplastic polymer made from styrene monomer; used in plastic model kits and plastic cutlery. Can be aerated to produce a foam (Styrofoam® ) and made into coffee cups or insulated food containers. Portland cement (12.1): Cement most commonly used in the production of concrete, consisting of a mixture of oxides of calcium, silicon, and aluminum. Positron (14.2): The antiparticle of an electron, having the same mass as an electron but with a charge of 11. Also called a b1 particle. Positron emission tomography (PET) (14.8): A medical imaging technique in which trace amounts of a positron-emitting isotope are injected into the body so that they accumulate in a particular organ. Detection of gamma rays produced in the annihilation of positrons is used to construct images of the organ being studied. Potential energy (9.2): Energy an object possesses because of its position. Precipitation reaction (3.3): Chemical reaction in which an insoluble solid (called the precipitate) is formed by the reaction of two or more solutions.

Precision (1.3): The spread in values obtained from a measurement. A precise measurement will produce similar results upon repeated observations. Preexponential factor (11.5): The preexponential term (A) in the Arrhenius equation. This term accounts for the frequency of collisions as well as any geometric or orientational constraints that influence the reaction rate. Also called the frequency factor. Pressure (P ) (5.2): The force exerted on an object per unit area. Primary battery (or primary cell) (9.8, 13.5): A single-use battery that cannot be recharged. Primary standard (5.1): Science-based limit established by the EPA for acceptable levels of criteria pollutants based solely on health considerations. (See also secondary standards.) (Also a substance of very high purity that can be used as a reference in determining the concentration of other substances.) Principal quantum number (n) (6.4): One of the three quantum numbers identifying an atomic orbital. Specifies the shell in which an orbital is found. Allowed values of n are positive integers. Products (3.2): Substances formed in a chemical reaction that appear on the right-hand side of a chemical equation. Protons (2.2): Positively charged subatomic particles found in an atom’s nucleus. Also H1 ions. PV-work (9.2): Work done by the expansion of a gas against an external pressure.

Q Quality factor (14.7): A variable to quantify the ability of a particular type of radiation to cause tissue damage. Used in calculating the equivalent dose. Also called the relative biological effectiveness. Quantum number (6.4): One of a number of indexes that identify a particular quantum state. An electron in an atom can be specified by a set of four quantum numbers.

R Radioactive decay (14.1): The spontaneous emission of subatomic particles or photons by an unstable nucleus, resulting in the formation of a new nuclide. Random error (1.3): Uncertainty associated with the limitations of the equipment with which a measurement is made; random error may cause the measurement to be either too high or too low, and so it can be minimized by averaging repeated measurements. Rate constant (k) (11.3): A proportionality constant relating the rate of a chemical reaction to the appropriate product of concentration terms for the species involved. Rate constants vary with temperature according to the Arrhenius equation. Rate law (11.3): A mathematical equation relating the rate of a chemical equation to the concentrations of reactants, products, or other species. Rate-determining step (11.6): The slowest step in a reaction mechanism. Reactant (3.2): Substance present at the outset of a chemical reaction that is consumed during the reaction; reactants appear on the left-hand side of a chemical equation. Reaction mechanism (11.6): Series of one or more elementary steps that provide a pathway for converting reactants into products.

Reaction quotient (Q) (12.3): Expression identical in form to the equilibrium constant but in which the concentrations do not correspond to equilibrium values. Comparison of the reaction quotient to the equilibrium constant predicts the direction of spontaneous change. Reaction rate (11.2): The change in the concentration of a reactant or product as a function of time, normalized for stoichiometry. Reactive intermediate (11.6): A chemical species that is generated in one step in a reaction mechanism and consumed in a later step. Intermediates do not appear in the overall chemical equation for a reaction. Reactor core (14.6): The central part of a nuclear reactor that contains the uranium fuel, moderator, coolant, and supporting structures. Redox reaction (13.2): A chemical reaction involving the transfer of electrons from one species to another. See also Oxidation–reduction reaction. Reduction (13.2): The gain of electrons by some chemical species in a chemical reaction. Reformulated gasoline (RFG) (4.6): Gasoline containing additives to provide at least 2% oxygen by weight. Reformulated gasoline is currently required in some urban areas in California, Texas, Kentucky, and Virginia, as well as in Washington, D.C. Refraction (6.2): The bending of light when it moves between two different media, such as air and the glass of a prism or lens. Relative biological effectiveness (RBE) (14.7): A variable to quantify the ability of a particular type of radiation to cause tissue damage. Used in calculating the equivalent dose. Also called the quality factor. Representative elements (2.5): Elements of Groups 1, 2, and 13–18, for which the highest lying valence electron occupies an s or p orbital. Also called main group elements. Resonance hybrid (7.5): The actual structure of a molecule represented by a combination of two or more different Lewis structures. Resonance structure (7.5): One of two or more individual Lewis structures that contribute to the resonance hybrid for a molecule. The resonance structures differ from one another only in the placement of electrons, but all show the same arrangement of the atoms themselves. Reversible (10.6): A process near equilibrium, which can be reversed by a small incremental change of some variable. Root-mean-square speed (5.6): The square root of the average of the squares of the velocities of a collection of particles.

S Sacrificial anode (13.8): A piece of a more reactive metal used to protect another metal from corrosion. The presence of the sacrificial anode assures that the part to be protected will be the cathode and hence will not be oxidized. Salt (3.3): Ionic compound that can be formed by the reaction of an acid with a base. Salt bridge (13.2): Device used to allow the flow of ions into the individual half-cell compartments in an electrochemical cell and maintain charge neutrality. Schrödinger equation (6.4): Mathematical equation for the allowed quantum mechanical wave functions for an atom or molecule.

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Scientific method (1.1): An approach to understanding that begins with the observation of nature, continues to hypothesis or model building in response to that observation, and ultimately includes further experiments that either bolster or refute the hypothesis. Scientific models (1.3): Empirical or mathematical descriptions that scientists create to make sense of a range of observations. Scientific notation (1.4): Representation of a number by a value between 1 and 10 times a power of 10. (Engineers in some fields prefer to use a number between 1 and 100 times a power of 10.) Scientific notation is a convenient way of displaying very large or small numbers, as well as a way of making clear which digits are significant. Sea-of-electrons model (8.3): Simplified description of metallic bonding in which the valence electrons of metal atoms are delocalized and move freely throughout the solid rather than being tied to any specific atom. Sea of instability (14.4): The region surrounding the band of stability in the chart of the nuclides, in which unstable nuclides are found. Second law of thermodynamics (10.4): The entropy of the universe is constantly increasing. For any spontaneous process, DSu > 0. Second-order reaction (11.4): Any reaction whose rate is given by a rate law of the form, Rate 5 k[A]2 or Rate 5 k[A][B], where A and B represent chemical species in the reaction (usually reactants). Secondary battery (or secondary cell) (9.8, 13.5): A battery that can be recharged by applying an external electric current to reverse the cell reaction and regenerate the reactants. Secondary quantum number (,) (6.4): One of the three quantum numbers identifying an atomic orbital. Specifies the subshell in which an orbital is found. Allowed values of , are related to the value of the principal quantum number (n). Secondary standard (5.1): Science-based limit established by the EPA for acceptable levels of criteria pollutants based on considerations of our environment and property. (See also primary standard.) Semimetals (2.5): Elements displaying some of the properties of metals and some of the properties of nonmetals. Shell (6.4): Set of orbitals in an atom that have the same value of the principal quantum number (n). Shielding (6.5): Reduction of the effective nuclear charge experienced by valence electrons due to the presence of inner shell electrons in an atom. Side reactions (4.4): Any reactions that lead to products other than those desired in an industrial or laboratory synthesis. Sigma (s) bond (7.6): A chemical bond in which the electron density is concentrated along the line of centers of the two atoms, such as occurs through the interaction of two s orbitals or the “end-to-end” interaction of two p orbitals. Significant figures (1.4): A method for expressing the degree of uncertainty in a measurement in which only those digits known accurately are reported. Simple cubic (sc) (8.2): Crystal lattice in which particles are arranged at the corners of a cube. Also called primitive cubic. Solid (1.2): State of matter in which a substance has a fixed shape and volume. Solubility (3.3): The ability of a compound to dissolve in solution. It can be expressed in terms of the maximum amount of solute for a given amount of solvent at a particular temperature.

562

Glossary

Solubility product constant (Ksp) (12.6): Equilibrium constant for the dissolution of a sparingly soluble compound, consisting of the concentrations of the resulting ions raised to appropriate stoichiometric powers. Soluble (3.3): Description of a solid that dissolves in solution to a reasonably large extent. Solute (3.3): The minor component(s) of a solution. Solution (3.3): A homogeneous mixture of two or more substances in a single phase. Solvent (3.3): The component present in the greatest amount in any solution. Sparingly soluble (12.6): Dissolving in water only to a very small extent in accordance with an appropriate solubility product constant. Specific heat (c) (9.4): The amount of energy required to raise the temperature of one gram of a substance by one degree (°C or K). Spectator ions (3.3): Ions that are present in a solution but do not actively participate in a chemical reaction. Spin paired (6.5): Used to describe two electrons, meaning they have opposite spins. To occupy the same orbital, two electrons must be spin paired. Spin quantum number (ms) (6.4): Quantum number specifying the spin of an electron. The two possible values (11/2 and 21/2) are usually referred to as “spin up” and “spin down.” Spontaneous process (10.2): Any thermodynamic process that takes place without continuous intervention. Note that being spontaneous does not necessarily mean that a process will be fast. Standard Gibbs free energy change (DG°) (10.7): Free energy change for a reaction in which all participating substances are present in their thermodynamic standard state (gases at 1 atm, solutions at 1 M, T 5 25°C). See also Gibbs free energy. Standard hydrogen electrode (SHE) (13.3): Electrode used to define the scale of standard reduction potentials, consisting of a platinum electrode at which 1 M H3O1 ions are reduced to H2 gas at 1 atm. Standard molar entropy (S°) (10.5): Absolute entropy of one mole of a substance at 25°C. Standard reduction potential (E°) (13.3): The potential of an electrochemical cell consisting of the given half-reaction and a standard hydrogen electrode; provides a convenient measure of the tendency of a species to undergo oxidation or reduction. Standard state (9.5): The most stable form of a substance at 25°C and 1 atm. Standard temperature and pressure (STP) (5.5): For gases, T5 0°C and P 5 1 atm. State function (9.6): A thermodynamic variable whose value depends only on the state of the system, not on its history. Statistical mechanics (10.3): A mathematical approach to thermodynamics which treats macroscopic systems by considering the statistical laws governing the interactions between their constituent particles. Stoichiometric coefficients (3.2): The numbers placed in front of each species to balance a chemical equation. Stoichiometry (3.2): The study of quantitative relationships between the amounts of reactants and products in a chemical reaction. Strong electrolyte (3.3): A substance that ionizes or dissociates completely in water to produce an aqueous solution in which only the constituent ions and not intact molecules are present.

Strong force (14.4): A force that acts at very short distances and is responsible for holding atomic nuclei together. Subshell (6.4): Set of orbitals in an atom that have the same values for the n and , quantum numbers. Superconductivity (8.7): Phenomenon in which a material can conduct an electric current without resistance; typically occurs only below some threshold temperature. Surface tension (8.5): Energy required to overcome the attractive forces between molecules at the surface of a liquid. Surroundings (9.3): The rest of the universe except for the thermodynamic system under consideration, with which the system can exchange energy. Symbolic perspective (1.2): Viewpoint of chemistry focusing on symbolic representations of the substances involved through formulas, equations, etc. Syndiotactic polymer (8.6): A polymer in which substituent groups are arranged on alternate sides of the polymer chain. System (9.3): The part of the universe that is being considered in a thermochemical problem. Systematic error (1.3): Uncertainty from some unknown bias or flaw in the equipment with which a measurement is made; a systematic error makes the measurement consistently either too high or too low and cannot be minimized by averaging repeated measurements.

T Temperature scales (1.4): Systems for measuring temperature, defined by choosing two reference points and setting a fixed number of degrees between them. Common temperature scales include Fahrenheit, Celsius, Kelvin, and Rankine. Termolecular (11.6): Description of an elementary step in which three reactant particles must collide with one another. Theoretical yield (4.4): The hypothetical maximum amount of product that can be obtained in a chemical reaction under ideal conditions. Thermochemical equation (9.5): A balanced chemical equation that includes the enthalpy change for the corresponding reaction. Thermochemistry (9.1): The study of the energetic consequences of chemistry. Thermocouple gauge (5.7): A sensor in which the conduction of heat away from a hot filament is used to determine pressure. Thermoplastic polymer (8.6): A polymeric material or plastic that can be repeatedly softened by heating and hardened by cooling. Thermosetting polymer (8.6): A polymeric material or plastic that forms covalent crosslinks such that it cannot be melted without decomposing. Third law of thermodynamics (10.5): The entropy of a perfect crystal of any pure substance approaches zero as the temperature approaches absolute zero. Titration (4.5): A process whereby a solution-phase reaction is carried out under controlled conditions using a known amount of one reactant so that the amount of the other reactant can be determined with high precision. Torr (5.2): A unit of pressure equal to 1 mm Hg. 760 torr 5 1 atm. Total ionic equation (3.3): A chemical equation in which all ions and molecules present in solution, including spectator ions, are shown.

Transition metals (2.5): The elements in Groups 3 through12 of the periodic table, in which valence electrons occupy d orbitals. Transmutation (14.6): Reaction in which one nucleus changes to another, either by natural decay or in response to some outside intervention, such as neutron bombardment. Triple bond (7.3): A covalent bond formed by the sharing of three pairs of electrons between the same pair of atoms. Turnover number (11.7): The effectiveness of a catalyst, expressed as the number of molecules that can react per catalyst binding site per unit time.

U UHMWPE (2.8): Ultrahigh molecular weight polyethylene. Polyethylene consisting of extremely long polymer chains and displaying unusually high strength. Uncertainty principle (6.4): It is impossible to determine both the position and momentum of an electron simultaneously and with arbitrary accuracy. Also called Heisenberg’s uncertainty principle. Uniform corrosion (13.1): Degradation of a metal surface by oxidation from uniform exposure to environmental moisture. Unimolecular (11.6): Description of an elementary step involving only a single reactant particle. Unit cell (8.2): The smallest collection of atoms or molecules that can be repeated in three dimensions to produce the overall structure of a solid. Units (1.4): Labels designating the type of quantity measured and the particular scale on which the measurement was made. Most measurements have no physical meaning if their units are omitted. Universal gas constant (R) (5.1): The proportionality constant in the ideal gas law; its value in SI units is 8.314 J K21 mol21, and a frequently useful value is 0.08206 L atm mol21 K21. Universe (9.3): In thermodynamics, the system under consideration plus its surroundings.

V Valence band (8.3): A band of orbitals or energy states populated by the valence electrons in the ground state of a bulk solid. Valence bond model (7.6): A description of chemical bonding in which all bonds result from overlap between atomic orbitals of the two atoms forming the bond. Valence electrons (6.5): The electrons in an atom’s highest occupied principal quantum shell, plus any electrons in partially filled subshells of lower principal quantum number. The electrons available for bond formation. van der Waals equation (5.6): An empirical equation of state used to describe quantitatively the behavior of a real gas by taking into account the volume of molecules and the interaction between them. Vapor pressure (8.5): The equilibrium gas phase pressure of a substance in equilibrium with its pure liquid or solid in a sealed container. Vaporization (9.5): The phase change in which a liquid is converted to a gas. Visible light (6.2): The portion of the electromagnetic spectrum that can be detected by human eyes, typically wavelengths between 400 and 700 nm. Volatile (8.5): Having a high vapor pressure; able to vaporize easily.

Glossary

563

Volatile organic chemical (VOC) (5.1): Carbon-based molecules that evaporate readily, are often emitted into the atmosphere, and are involved in the formation of smog. VSEPR theory (7.8): Valence-shell electron-pair repulsion theory. A simple method for predicting the shapes of molecules; it states that molecules assume a shape that allows them to minimize the repulsions between bonds and lone pairs in the valence shell of the central atom.

W Wave function (c) (6.4): A mathematical function providing the quantum mechanical wave representation of an electron in an atom; a solution to the Schrödinger wave equation. Wavelength (l) (6.2): The distance between adjacent corresponding points (peaks, valleys, etc.) of a wave. Wave-particle duality (6.2): The notion that, in some situations, light or electrons are best described as waves, while in other cases a particle description works better. Weak electrolyte (3.3): A substance that ionizes or dissociates only partially in water to produce an aqueous solution in which both individual ions and intact molecules are present.

564

Glossary

Work (w) (9.2): A transfer of energy accomplished by a force moving a mass some distance against resistance.

Y Yield strength (1.6): A measure of the amount of force required to produce a specified deformation of a material.

Z Zero-order reaction (11.4): Any reaction whose rate is independent of the concentrations of all participating substances. Zinc-air battery (13.5): A primary cell in which zinc is oxidized at the anode and oxygen is reduced at the cathode. Because oxygen can be obtained on demand from the air, these batteries can have high energy density and are suitable for use as emergency power supplies for cell phones or other devices.

Index Italicized page numbers indicate illustrations, page numbers followed by “t” indicate tables, those followed by “q” indicate end-of-chapter questions, and those followed by “b” indicate glossary entries.

A ABS, 270 Absolute temperature, 142 Absolute temperature scale, 135b Absolute zero defined, 132, 135 third law of thermodynamics, 327–329, 328t, 328, 342q Absorbed dose, 498b, 498t Acceptor level, 254b Accreditation Board for Engineering and Technology, 2 Accuracy, 10, 10b Acetates, 74t Acetic acid acid ionization constant, 422t, 423–424 characteristics of, 78t as weak electrolyte, 76 potassium hydroxide reaction with, 79–80 Acetone, 263t Acetylene, 343 Acid(s), 419–425, 433q. See also Acid–base reactions acid ionization constant, 422b, 422t Arrhenius, 420 BrØnsted–Lowry theory, 420–421 chemical equilibrium, 419–425 conjugate acid–base pairs, 421, 421b, 433 defined, 78b as electrolytes, 78 strong, 78, 78t titration, 115b, 115–116, 116 weak, 78, 78t, 421–425 Acid anhydride, 424 Acid ionization constant, 422b, 422t Acid spills, 80 Acid–base reactions, 78–81 acid ionization constant, 422b, 422t BrØnsted–Lowry theory of acids and bases, 420–421 in nonaqueous medium, 421 titration, 115b, 115–116, 116 weak acids and bases, 421–425 Acrylic acid, 433 Acrylonitrile, 270 Actinides, 46b Activated complex, 368b Activation energy, 367b, 372 Active electrolysis, 458b, 460–461 Actual yield, 113b Addition, 17 Addition polymerization, 266, 266, 267 Addition polymers, 266, 266–268, 267 Addition reactions, 52 Additives to gasoline, 110–111, 117 to paint, 465

to polymers, 272, 427–428 Adhesion, 265b Adiabatic processes, 322 Adipic acid, 97 Admixtures, 392b, 393t ADN. See Ammonium dinitramide Advanced recycling, 327 Air, 125, 126t. See also Atmosphere Air bags, 95 Air entraining admixtures, 393 Air pollution, 126–129 fuel additives, 117 gas pressure and, 132 questions regarding, 152q volcanoes, 137–138 Alchemy, 489 Alcohol thermometer, 15 Alcohols, 52, 52t Aldehydes, 53t Alkali metals, 45b Alkaline batteries, 309, 309, 310t, 315, 453b, 453–455, 454, 471 Alkaline earths, 45b Alkanes, 100b in gasoline, 100 list, 101t molecular structure of, 100, 101t, 102 Alkenes, 52t Alkylperoxy radical, 380 Alkynes, 52t Alloys, 256 Alpha decay, 477b, 477–478 Alpha particle, 477b, 496 Alpha rays, 476b, 477 Alternating copolymers, 270b, 270 Alternative fuels, 117–118, 122q Alumina, 9, 23 Aluminum, 2–4 bicycle frames composed of, 24, 24t, 25 chemical properties, 5–6 commercial production, 123 corrosion, 438 ionization, 203 ionization energy, 190t physical properties, 5–6, 24, 24t questions regarding, 27q refining of, 11, 23, 458–460, 459 smelting of, 23, 23–24 in solid fuel booster rockets, 111–112 specific heat, 290t sulfuric acid reaction with, 124 uses of, 5, 26 Aluminum alloys, 25 Aluminum cans, 2–5, 9, 11, 291, 335 Aluminum composites, 25 Aluminum hydroxide, 120–121 Aluminum ore

description of, 4, 8, 11 digestion, 23, 23 processing, 22–24, 23 Aluminum oxide, 9, 9 Amides, 52t Amines, 52t Ammonia in Andrussow process, 403, 409 description of, 26, 78, 78t, 277 formation, 109, 403 Haber process for, 383 reaction with oxygen, 383 uses of, 109 van der Waals constant, 146t Ammonium dinitramide, 155 Ammonium ion, 55t Ammonium nitrate, 77, 86, 120 Ammonium perchlorate, 111–112 Ammonium sulfate, 122 Amorphous structures, 243b Ampere (unit), 14t, 461 Amphiprotic, 420 Amphoteric, 420b Amplitude, 161b, 162, 162 Amu. See Atomic mass unit Andrussow process, 403, 409 Anions, 38, 38t, 38b common anions, 55t formation, 204–207, 205–206, 235 Anode, 441b, 448 Antibonding molecular orbital, 251b Antineutrino, 478 Aqueous reactions, 77 Aqueous solutions, 72–81, 93, 94q–95q in chemical equations, 67, 76 defined, 72b dilution formula, 90, 116 electrolytes, 75, 76 solubility guidelines for ionic compounds, 74t Argon electron configuration, 185 ionization energy, 190t van der Waals constant, 146t Argon gas, 185 Arrhenius, Svante, 420 Arrhenius acids and bases, 420 Arrhenius equation, 368b, 368–369, 369, 371, 373 Arsenic, 120, 254–255 Aspdin, Joseph, 392 Aspirin, 385 Atactic polymers, 268b -ate (suffix), 55 Atmosphere nuclear reactions in, 475 ozone depletion, 348–350, 349, 365, 365–366 ozone in, 348–350, 349, 360

565

Atmosphere (unit), 131b Atmospheric pressure, 129, 129, 130 Atom(s), 6b. See also Atomic structure Bohr, 172, 172–173 components of, 33, 33–34 isotopes, 34b, 34–35 mass number, 34, 34b quantum mechanical model, 173–181, 196q, 256 word origin, 6 Atomic absorption spectroscopy, 199 Atomic mass, 36–37, 43, 60q Atomic mass unit, 34b, 36, 36 Atomic number, 34, 34b, 187, 188, 478 Atomic orbitals. See Orbital(s) Atomic properties, 187–192, 197q Atomic radius, 187, 188 Atomic size, 187, 187–188 Atomic spectra, 169b, 169–173, 196q Atomic structure, 33–37, 59, 60q. See also Orbital(s) basic elements, 33, 33–34 Bohr atom, 172, 172–173 quantum mechanical model, 173–181, 196q, 256 Atomic symbol, 35–36, 36t, 39 Atomic theory, 58 Atto-, 14t Aufbau principle, 183b, 183–184, 186, 251 Automobiles. See also Gasoline catalytic converters, 97, 117, 122, 377, 378, 388 exhaust, 95, 102–103, 119 Average rate of reaction, 352–353, 353 Average speed, 144, 144b Avogadro’s law, 132b Avogadro’s number, 82b, 82–83, 93

B Balanced chemical equations, 68–72, 72, 104–108, 118 “Ball and stick” model, 48, 229t–230t Band energy diagrams, 252, 252 Band gap, 252b Band of stability, 485b, 485, 500 Band theory, 250b, 251–252, 275 Barium carbonate, 80–81 Barium chloride, 80–81 Barium hydroxide, 78t Barium sulfate, 75, 433 Barometer, 130, 130b, 131 Barrel plating, 460, 460 Base(s), 419–425, 433q. See also Acid–base reactions Arrhenius acid, 420 base ionization constant, 422b BrØnsted–Lowry theory, 420–421 chemical equilibrium, 419–425 conjugate acid–base pairs, 421, 421b defined, 78b as electrolytes, 78 strong, 78, 78t titration, 115b, 115–116, 116 weak, 78, 78t, 421–425 Base ionization constant, 422b Batteries, 453b alkaline, 309, 309, 310t, 315, 453b, 453–455, 454, 471

566

Index

description of, 308–311, 309, 310t, 315q, 453–458, 467 dry cell, 454 fuel cells vs., 457 lead-acid storage, 456, 456b, 457 limitations, 457–458 lithium, 310t lithium ion, 310t mercury, 454, 454b, 454 nickel-cadmium, 310t, 445b, 455, 455 nickel-metal-hydride, 310t, 456, 456b, 456 primary cells, 309b, 309–310, 310t, 453b, 453–455 questions regarding, 471q secondary, 309b, 455b, 455–456 zinc-air, 310t, 454b, 454–455, 455, 471 Bauxite, 4, 8, 8, 22–23, 22–23 bcc structure. See Body-centered cubic structure Becquerel, 481b Bent molecule, 229t Benzene, 221, 221 Benzopyrene, 124 Beryllium, 184, 190t Beta decay, 478b, 478–479 Beta particle, 477b, 496 Beta rays, 476b, 477 Beverage cans. See Aluminum cans BHET. See Bis(2-hydroxyethyl) terephthalate Bicycle frames, 24t, 24–25, 29q Bimolecular steps, 374b Binary compounds, 53, 53b, 54 Binding energy, 168, 168b, 487b, 487–488, 488, 500 Biocompatibility, 201b Biomedical engineering, 201–202 Bis(2-hydroxyethyl) terephthalate, 319, 319 Bismuth selenide, 124 “Black lights,” 165 Block copolymer, 270b, 270 Body-centered cubic structure, 245, 245, 247 Bohr, Niels, 172 Bohr atom, 172, 172–173 Bohr model, 172b Bohr radius, 179 Boiling of water, 7, 7 Boiling point, 15, 15, 263t, 263–264, 264 Boltzmann distribution, 262, 367 Bomb calorimeter, 293, 293 Bond. See Chemical bonds Bond energy, 207b Bond length, 207b Bond polarity, 213, 213–214, 214, 237q Bonding molecular orbital, 251b Bonding pair, 211b Booster rockets, 111–112 Borates, 427–428 Borax, 427–428 Boric acid, 427–428 Boron, 58 electron configuration, 184 ionization energy, 190t isotopes, 58–59 Boron neutron capture therapy, 505q Borosilicate glass, 428 Bose-Einstein condensates, 5 Bottle recycling, 320, 339 Boundary, 286b Boyle, Robert, 132

Boyle’s law, 132b, 152 Branching, 57 British thermal unit (unit), 282 Brfnsted, Johannes, 420 Brfnsted–Lowry theory of acids and bases, 420–421 Bromide ion, 55, 55t Bromides, 74t Bromine boiling point, 263t chemical reactivity, 45 Bromine pentafluoride, 231–232 Btu (unit), 282 Buckminsterfullerene, 241, 273 Bulletproof vests, 58 Burning, 5. See also Combustion Butane, 101t Butanoic acid, 422t

C Calcium carbonate, 75, 97, 123 Calcium fluoride, 416 Calcium hydroxide, 78t, 403 Calcium ion, 55t Calcium phosphate, 201, 417–418 Calibration, in calorimetry, 293 Calorie, 286b Calorimetry, 289b, 293, 293–294, 313q Candela (unit), 14t Candle wax, 382–383 Capacitance manometer, 148, 148–149, 150 Capacitor, 148 Carbon. See also Organic compounds atomic mass, 36, 36 atomic number, 34, 36 atomic symbol, 36 covalent bonding, 50 crystal structure, 247–248, 248–249 diamond, 241, 242, 247–248, 248, 252 double bond, 211 fullerenes, 241, 273 graphite, 242, 242, 247, 249, 249, 260–261, 261 ionization energy, 190t isotopes, 35–36 Lewis symbols, 210 mass spectrum, 34–35, 35 molar mass, 82 nanotubes, 242, 292 phase diagram, 242, 242 questions regarding, 275q–276q single bond, 211 triple bond, 211 Carbon-12, 36, 485–486 Carbon-13, 36, 485 Carbon-14 description of, 475–476, 478 formation of, 483 half-life of, 482–483, 484t radiocarbon dating uses of, 483–484 Carbon atom, 50 Carbon dating, 475–476, 501q Carbon dichlorodifluoride, 263t Carbon dioxide chemical bonds, 43 dissolving in water, 412 dry ice, 21–22 mass, 135–136

Maxwell-Boltzmann distribution, 143, 143 removing from industrial exhaust, 140–141 van der Waals constant, 146t Carbon fiber composites, 25 Carbon tetrachloride, 263t, 277 Carbonate anion, 55t Carbonates, 74t, 75 Carbonation, 394 Carbonic acid, 422t, 432 Carboxylic acids, 52t Carnot cycle, 322, 322 Catalysis, 376–379, 388q Catalysts, 41 chemical equilibrium and, 415 heterogeneous, 377, 377b homogeneous, 377b rates of reaction, 376–379, 388q Ziegler-Natta, 268 Catalytic converters, 97, 117, 122, 377, 378, 388 Cathode, 441b, 448 Cathodic corrosion protection, 466, 466b, 466 Cations, 38, 38t, 38b common, 55t formation, 202–203, 235 Caustic soda, 23 ccp structures. See Cubic close-packed structures Cell adhesion, 202 Cell notation, 441b Cell potential, 442b description of, 444–453, 467 equilibrium and, 450–453, 452 equilibrium constants and, 452, 452–453 free energy and, 450–451 measuring of, 444, 444–445, 446 nonstandard conditions, 449, 467 questions regarding, 468q–471q Celsius scale, 15, 15 Centi-, 14t CFCs. See Chlorofluorocarbons Chain reaction, 492b Chain termination, 267 Chapman, Sydney, 348 Chapman cycle, 348–350, 373, 376–377 Charge, 34, 461 Charles, Jacques, 132 Charles’s law, 132b, 152 Chart of the nuclides, 485b, 485 Chemical bonds, 40b bond energy, 207 bond length, 207 bond polarity, 213, 213–214, 214, 237q chemical reactions, 209 covalent bonds, 43b, 47, 207–211, 235, 236q–237q description of, 6, 42–43, 59, 235, 239, 275 double bond, 211, 232 electronegativity, 212, 212b, 237q energy and, 207–209, 208, 208–209, 299–300 intermolecular forces, 256b, 256–261, 257–258, 275, 276q–277q ionic bonds, 42, 42, 202–207, 235, 236q Lewis structures, 210–211, 215–221 metallic bonding, 43, 43b, 43, 249–251, 250–251, 276q molecular structure and, 209–211 notation, 48, 49–50 octet rule, 210–211 orbital overlap, 221–224, 222–223, 238q

peptide bonds, 47 pi bonds, 223, 223 polar bond, 213 questions regarding, 61q resonance, 220b, 220–221, 221, 273–274 sigma bonds, 223, 223 single bond, 211 in solids, 249–256, 275 triple bond, 211 valence bond model, 222 valence electrons and, 185 Chemical calculations, 93q–94q. See also Stoichiometry Avogadro’s number, 82b, 82–83, 93 dilution, 90 empirical and molecular formulas, 86–88 molar mass, 82, 82b, 84–90, 93, 95q–96q, 105, 107 molarity, 88–89 mole ratios, 83–84, 104b, 104–108, 118 moles, 14t, 82b, 82, 82–85, 87, 93, 95q–96q ratios, 19–21, 26, 104–108, 106, 118 STP conditions, 140–141 Chemical changes, 7 Chemical energy, 284b Chemical equations, 67b, 93, 93q–94q, 103. See also Chemical equilibrium; Chemical reactions; Stoichiometry aqueous reactions, 77 balancing of, 68–72, 72 interpreting, 81 mole ratio, 83–84, 104b, 104–108, 118 molecular equation, 77b net ionic equation, 77b reversing, 402–403 total ionic equation, 77b two-way arrow, 78–79 unbalanced, 68 writing of, 67–68 Chemical equilibrium, 391–435, 429q–430q acids and bases, 419–425 catalysts and, 41, 376–379, 388q, 415 common ion effect, 418b, 418, 418–419 concentration changes and, 410–412, 411t dynamic equilibrium, 263b, 396b, 429 electrochemistry, 450–453 equilibrium concentration, 405–410, 431q–432q equilibrium constants, 398–405, 429, 430q–431q equilibrium expression, 398b, 401–402 forward reactions, 394–396, 395–396, 429 free energy and, 425–427, 426 gas phase equilibria, 399–400 heterogeneous equilibrium, 400b homogeneous equilibrium, 400b law of mass action, 398 LeChatelier’s principle, 410b, 410–415, 429, 432q molar solubility, 416b, 416–418 pressure changes and, 412–413, 413 rate laws and, 397–398 reverse reactions, 394–396, 395–396, 429 reversing the chemical equation, 402–403 solubility equilibria, 415–419, 432q–433q solubility product constant, 415–416, 416t temperature changes and, 414–415, 415t Chemical explosions. See Explosions

Chemical formulas, 40b description of, 40–41, 85–86 questions regarding, 93q–94q rules for, 41 use in writing chemical equations, 67 Chemical kinetics, 347–390 activation energy, 367b, 372 Arrhenius behavior, 368b, 368–373, 369, 370, 373 catalysis, 41, 376–379, 388q, 415 differential rate law, 353–358, 354b, 382 half-life, 364–366 integrated rate laws, 358–366, 382, 385q–387q order of the reaction, 354b, 354–355 ozone depletion, 348–350, 349 rate constant, 354b, 354–355, 383–384 rate-determining step, 376, 376b reaction mechanisms, 373–376, 387q–388q reaction rates, 350–356, 382, 382q–383q temperature and, 366–373, 367–369, 371, 372, 387q Chemical nomenclature, 53–56, 59, 62q binary compounds, 53, 54 covalent compounds, 53–54 ionic compounds, 54–56, 55t Chemical properties, 5b Chemical reactions, 64, 93, 382. See also Chemical equations; Chemical equilibrium; Chemical kinetics; Thermodynamics acid–base reactions, 78–81 activated complex, 368b activation energy, 367b, 372 addition polymerization, 266, 266, 267 addition reactions, 52 aqueous solutions, 72–81 Arrhenius behavior, 368b, 368–373, 369, 370, 373 Avogadro’s number, 82b, 82–83 catalysis, 41, 376–379, 388q, 415 chemical bonds, 209 condensation polymerization, 268, 269 depolymerization, 327, 333 efficiency, 113–114 endothermic, 296b, 299–300, 305, 322, 414–415 exothermic, 296b, 299–300, 305, 322, 340, 414 explosions, 64, 65–66, 66 formation reactions, 300b, 303–305 forward and reverse reactions, 394–396, 395–396, 429 free energy and, 333–335, 339, 343q half-reactions, 439b, 446, 452, 467 heat and, 299–300 heat of reaction, 299b, 300, 314q irreversible, 333b limiting reactants, 108b, 108–112, 120q–121q measuring progress of, 14 mole, 82b, 82, 82–84 neutralization, 79b order of, 354b, 354–355 oxidation–reduction, 308b, 438–440, 439, 467 percentage yields, 113–114, 121q photochemical reaction, 68b, 127b polymerization. See Polymerization precipitation, 80b, 80–81 rate-determining step, 376, 376b rates of reaction, 350–356, 382, 382q–383q

Index

567

Chemical reactions (cont. ) reaction mechanisms, 373–376, 387q–388q reactive intermediate, 374b reversible, 333b reversing of, 402–403 side reactions, 113b, 114 spontaneity, 320–322, 339, 339q–340q stoichiometry, 69b, 99–124, 403–404 STP conditions, 140–141 theoretical yields, 113–114, 121q titration, 115b, 115–116, 116 Chemistry, 27q–29q as “central science,” 4 conceptual problems, 21–22 as empirical science, 9, 26 green, 91b, 91–92, 96q, 113 macroscopic perspective, 4b, 4–6, 81 microscopic (particulate) perspective, 4b, 6–7, 81 models, 11b, 11–12, 26 numbers and measurements, 12–18 observations, 9–11, 26 symbolic representation, 8–9, 9 visualization in, 22–24 Chlorate ion, 55t Chlorates, 74t, 75 Chloride ion, 38t, 40, 55 Chlorides, 74t Chlorine, 48 chemical reactivity, 45 ionization energy, 190t isotopes, 37, 62 oxyanions, 55t Chlorine gas, 40, 408–409 Chlorite ion, 55t Chloroacetic acid, 422t Chlorofluorocarbons, 124, 156, 216–217, 366, 377 Chlorophyll, 98 Chlorotrifluoromethane, 263t Chromium, 124, 451 Clean Air Act, 98 Close-packed structures, 244, 244, 247 Coatings, for corrosion prevention, 465 Coherent light, 193b Cohesion, 265b Cold vulcanization, 123 Column (periodic table), 44 Combustion, 5 efficiency, 114 incomplete, of alkanes, 102, 103, 343 limiting reactant, 109 matches, 110 of acetylene, 343 of ethane, 124 of gasoline, 102, 107, 122, 287, 321 of hydrocarbons, 68, 68, 69–70, 100, 102, 104, 209, 285 of hydrogen, 108 of magnesium, 154 of methane, 68, 104–105, 299–300, 302, 302 of naphthalene, 294 of octane, 102, 103, 120 of propane, 69–70, 119, 304, 387 percentage yield, 114 reaction rate, 350, 353 Combustion analysis, 124

568

Index

Common ion effect, 418b, 418, 418–419 Compound(s), 40, 59, 61q. See also Chemical nomenclature; Molecules binary, 53, 53b, 54 chemical bonding, 42–43 fluorine. See Fluorine compounds ionic. See Ionic compounds Compound nucleus, 489b Compression, 24 Computer chips, 496 Concentrated solutions, 73, 93 Concentration chemical equilibrium and concentration changes, 410–412, 411t dilution, 90 molarity, 88–89 of solutions, 73b rate laws, 353–366, 383q–385q Conceptual problems, 21–22 Concrete, 392–394, 393, 398, 424 Condensation polymerization, 268, 269 Condensation polymers, 268b, 268–270 Condensed phases, 243–249, 276q. See also Liquid(s); Solid(s) Conduction, 75, 76 Conduction band, 252b, 253 Conductivity, 251–252, 275 Conjugate acid, 421b Conjugate acid–base pairs, 421, 421b, 433 Conjugate base, 421b Conservation of energy, 286–287, 312, 313q Conservation of matter, 68 Control rods, 492b, 493 –COOH functional group, 78 Coordination number, 247, 247b Copolymers, 270b, 270, 270–271 Copper, 60q electrochemical cell, 447–448 galvanic cell, 440, 440–444, 441b, 442, 446, 467, 468q, 470 nitric acid reaction with, 121 silver nitrate reaction with, 438–439, 439 specific heat, 290t sulfuric acid reaction with, 120 Copper(I) ion, 55t Copper(II) ion, 55t Copper(II) oxide, 121 Core electrons, 185b Corrosion, 437b description of, 5, 437–438, 450, 467 galvanic, 437b, 442–444, 448, 451, 465 illustration of, 437 prevention of, 465–467, 472q–473q questions regarding, 468q space program, 437, 466 uniform, 437b, 437 Cosmic rays, 475b, 475–476, 501q Coulomb potential, 175 Coulomb (unit), 34, 461 Coulombic force, 39, 39, 42 Coulomb’s law, 38b, 38–39, 59, 204, 486 Covalent bonds, 43b, 207b description of, 207–211, 235 octet rule, 210–211 questions regarding, 236q–237q in silicon-based polymers, 47

Covalent molecules chemical nomenclature, 53–54 orbital hybridization, 226 Criteria pollutants, 127b Critical mass, 492b Cross-linking of polymers, 272 Cryolite, 23, 123, 459 Crystal structure coordination number, 247, 247b packing efficiency, 243–245, 245–246 periodic table and, 247–248, 248 third law of thermodynamics, 327–329, 328t, 328, 342q Crystalline solids, 243–249, 245–246, 248 Crystalline structures, 243b Cubic close-packed structures, 245, 245, 247 Cumene, 97 Curie, 481b Curl, Robert, 273 Cyanide anion, 55t

D Dacron, 268, 269 Dalton, John, 137 Dalton (unit), 34 Dalton’s law of partial pressures, 137 de Molyens, Frederick, 192 Deca-, 53t Decaborane, 118 Decane, 101t Decay positron, 480 radioactive. See Radioactive decay Decay constant, 481b Decay series, 486b Deci-, 14t Deductive reasoning, 11b, 26 Degree Celsius (unit), 15 Degree Fahrenheit (unit), 15 Degree of polymerization, 267b Degree Rankine (unit), 16, 135 Deka-, 14t Delta, 287 Dendrochronology, 483 Density, 5b, 6, 6 Depolymerization, 327, 333 Derived units, 14 Detonation, 65 Deuterium, 494 Di-, 53t Diamond combustion of, 321 crystal structure of, 247–248, 248 description of, 241, 252 heat of formation, 303 industrial uses of, 241 manufacture of, 241 phase change, 314 Diborane, 94 Dice, probability, 323, 323 Dichlorodifluoromethane boiling point of, 263t Lewis structure of, 216–217 Diethyl ether, 344 Diethylaluminum chloride, 41 Differential rate law, 353–358, 354b, 382 Diffraction, 174b

Digestion (aluminum ore), 23, 23 Dilute solutions, 73 Dilution, 90, 90b, 93 Dilution formula, 90, 116 Dimensional analysis, 19b, 26 2, 2-Dimethylpropane, 100 Dinitrogen oxide, 54, 367 Dinitrogen pentoxide, 54, 54, 356, 375–376, 385 Dinitrogen tetroxide, 54 Dinitrogen trioxide, 54 Dipole, 213b Dipole–dipole forces, 258, 258, 275 Dipole–dipole interactions, 258 Discrete spectra, 170 Dispersion forces, 256b, 256–258, 257, 259, 275 Dissociation, 76 of hydrochloric acid, 78 of sodium hydroxide, 78 of strong acid, 78 of strong base, 78 Dissociation reaction, 76 Dissolving, 73. See also Solutions Distance units, 14t Distillation, 327 Distribution function, 142b, 142–143 Division, 17 Donor level, 253b Doping, semiconductors, 253, 253b, 253, 273–274 Dosimeter, 497b, 497 Double bond, 211b, 232 Drug delivery, molecular scale engineering for, 234–235, 238q Dry air, 154 Dry cell, 454 Dry ice, 21–22 Ductile, 249b Dynamic equilibrium, 263b, 396b, 429

E E 5 mc 2, 487, 500 Economics of plastics recycling, 335–338, 337t, 343q–344q Edison, Thomas, 453 Effective nuclear charge, 182b, 188 Efficiency of chemical reaction, 113b, 114 Einstein, Albert, 167 Elastic modulus, 24b, 24t Elasticity, of polymers, 272 Electric charge, 34, 38–40, 59 Electric current charge, 461 units of, 14t Electric lights, 158, 159–161, 160 Electrical conductivity, 251–252, 275 Electrical energy, 285 Electricity power plants, 298, 299 Electrochemical series, 447 Electrochemistry, 436–473, 441b. See also Oxidation–reduction reactions batteries. See Batteries cell potential. See Cell potential chemical equilibrium and, 450–453 corrosion. See Corrosion electrolysis. See Electrolysis electromotive force, 442b

equilibrium constants, 452–453 Faraday constant, 449 galvanic cell, 440, 440–444, 441b, 442, 446, 467, 468q, 470 galvanic corrosion. See Galvanic corrosion half-reactions, 439b, 446, 452, 467 Nernst equation, 449–450, 467, 469 oxidation–reduction reactions, 308b, 438–440, 439, 467 standard hydrogen electrode, 444–445, 445, 468–469 standard reduction potential, 445–449, 446t, 446b, 447, 468, 471 Electrodes, 441b Electrolysis, 7, 7, 458b active, 458b, 460–461 aluminum refining, 458–460, 459 description of, 458–464, 467 electroplating, 460b, 460, 460–462, 465 masses of substances used in calculations, 463–464 passive, 458b, 458–460 polarity and, 458, 459 questions regarding, 471q–472q stoichiometry and, 461–464, 472q Electrolytes, 75b, 76 Electrolytic cell, 458 Electromagnetic spectrum, 161, 161–169, 164, 195q–196q Electromotive force, 442b Electron(s), 33b binding energy, 168 in Bohr model, 172 bonding pair, 211 chemical bonds, 42, 42–43 configuration. See Electron configuration core electrons, 185b excited state, 172b lone pair, 211 mass, 34 in metallic bonding, 43, 43 orbitals, 173b, 178–181, 179–181, 182 Pauli exclusion principle, 182b, 197q in quantum model of atom, 173–174 spin paired, 182 spin quantum number, 181 valence electrons, 185b Electron affinity, 191b, 191, 191–192, 202 Electron capture, 479b, 479–480 Electron configuration, 183b, 184 description of, 185, 195 orbitals, 182–183, 183 periodic table and, 185–187, 186, 197q questions regarding, 197q Electron diffraction, 174 Electron sea model of metallic bonding, 43, 250, 275–276 Electron volt (unit), 475b Electronegativity, 212, 212b, 237q, 249 Electroplating, 460b, 460, 460–462, 465 Element(s), 6b. See also Periodic table; specific element atomic symbols, 35–36, 36t isotopes, 34b, 34–35 Elemental analysis, 86b, 86–88 Elementary steps, 374b EMF. See Electromotive force

Empirical equation, 146 Empirical formulas, 40b, 86 Endothermic reactions, 296b, 299–300, 305, 322, 414–415 Energy, 280–312, 312q–313q. See also Heat activation, 367b batteries. See Batteries binding, 487b, 487–488, 488, 500 chemical bonds and, 207–209, 208, 208–209, 299–300 conservation, 286–287, 312, 313q defined, 284 electricity production, 298 explosions, 66 first law of thermodynamics, 287, 312 forms, 284–285 fuels, 307 heat and work, 285 kinetic. See Kinetic energy of photon, 167 potential. See Potential energy second law of thermodynamics, 326–327, 339, 341q stoichiometry and, 305–307, 314q thermochemical equation, 300b transformation, 286, 313q units, 285–286 waste, 288t, 288–289, 289 Energy conversion, 288t, 288–289 Energy density, 294, 307, 307b Energy economy, 280b Energy use energy sources, 282 questions regarding, 312q U.S. production and consumption, 282, 282–283, 283 world economy and, 281–284, 281–284 Energy-level diagram, 171 Enthalpy, 295–301 definition of, 295–296 Hess’s law, 301–305, 314q, 333 questions regarding, 314q spontaneity and, 321–322, 339 Enthalpy diagram, 302, 302 Enthalpy-driven processes, 331 Entropy, 322b, 322–326 changes, in processes, 324–326 definition, 324 description of, 322–324, 328, 339 questions regarding, 340q–341q third law of thermodynamics, 327–329, 328t, 328, 342q Entropy-driven processes, 331 Environmental chemistry, 91 Environmental issues. See also Air pollution; Pollutants CFCs. See Chlorofluorocarbons Montreal Protocol, 377 ozone depletion, 348–350, 349 smog, 125, 127, 379–380, 387 Environmental Protection Agency, 91 Equations Arrhenius, 368b, 368–373, 369, 371, 373 chemical. See Chemical equations empirical, 146 molecular, 77, 93 net ionic, 78, 81, 93, 94q–95q

Index

569

Equations (cont. ) nuclear, 500 Schrödinger, 174–175, 177 thermochemical, 300, 305, 414 total ionic, 77 van der Waals, 146 Equilibrium. See Chemical equilibrium Equilibrium concentration, 405–410, 431q–432q Equilibrium constants, 398–405, 429, 430q–431q, 452, 452–453, 470–471 Equilibrium expression, 398b, 401–402 Equivalent dose, 498b, 498t Error, measurement, 10, 12 Ethane, 101t boiling point, 263t combustion, 124 formation, 119 molecular structure of, 101t Ethanol, 52 as fuel additive, 94, 117, 305 heat of formation, 305 Ethers, 52t Ethylene chemical formula of, 40, 40 molecular formula of, 56 polymerization of, 56–57 structural formula of, 56, 57 Ethylenedinitramine, 85 Europium, 276 Exa-, 14t Excited state, 172b Exothermic reactions, 296b, 299–300, 305, 322, 340, 414 Explosions, 64, 65–66, 66, 91 destructive force, 66, 66 grain elevator, 385 Explosives, 77, 83–86, 91–92, 93q, 96q, 98, 120 Exposure, radiation, 498, 498t

F Face-centered cubic structure, 245–247 Factor-label method, 19b, 26 Fahrenheit scale, 15, 15 Faraday, Michael, 449 Faraday constant, 449b fcc structure. See Face-centered cubic structure Feedstock recycling, 327 Femto-, 14t Femtosecond (unit), 14 Fermi level, 252 Ferric, 55 Ferrous, 55 Filament, lightbulb, 159–160, 160, 186, 192 Film-badge dosimeter, 497, 497 Fire retardants, 428 Firearms, 91 Fireworks, 199 First ionization energy, 188, 190 First law of thermodynamics, 287, 312 First-order reactions half-life, 365 integrated rate law, 360–361 radioactive decay, 365 Fission, 490, 490–492, 492, 504q Flash electroplating, 461–462 Fluorescent lights, 160, 160–161, 169–170, 173, 185, 195q, 197q–198q

570

Index

Fluoride ion, 204 Fluorides, 74t Fluorine, 212 atoms, 211 chemical reactivity, 45 hydrogen reaction with, 213, 213 ionization energy, 190t Fluorine compounds bond polarity, 214 chlorofluorocarbons, 124, 156, 216–217, 366, 377 Freon-12, 124, 216–217 PTFE, 209, 209t Teflon, 209 Fluorocarbons, 209t Flux, 459 Fly ash, 392–393 Formaldehyde in auto exhaust, 119 boiling point, 263t from incomplete combustion of propane, 70–72 Formation reactions, 300b, 303–305 Formic acid, 422t Formula unit, 42b Forward reactions, 394–396, 429 Free energy, 339 cell potential and, 450–451 chemical equilibrium and, 425–427, 426 chemical reactions and, 333–335, 343q Gibbs, 330–333, 339, 342q–343q Helmholtz, 330, 342 nonstandard conditions, 426–427 spontaneous change and, 330–333 standard Gibbs, 333b work and, 333 Free radicals, 57, 57b Freezing point of water, 15, 15 Freon-12, 124, 216–217, 366 Frequency, 162b, 162 Frequency factor, 368b Fuel(s). See also Gasoline alternative, 117, 122q description of, 89, 91, 100–103 energy density, 307, 307b, 308t hydrocarbons, 68, 68, 69–70 hydrogen as, 106 oxygenated, 117 percentage yield, 114 questions regarding, 119q rocket propellant, 94, 111–112, 117 solid fuel booster rockets, 111–112, 155 Fuel additives, 110–111, 117–118, 122q Fuel cells, 457, 457b Fuel rods, 492b, 493, 493 Fuller, Buckminster, 241 Fullerenes, 241, 273 Functional groups, 52, 52b, 52t–53t Fusion, 494–495, 504q

G Gadolinium-153, 505q Gallium, 60, 254–255, 344 Gallium arsenide, 256 Galvani, Luigi, 441 Galvanic cell, 440, 440–444, 441b, 442, 444, 446, 467, 468q, 470

Galvanic corrosion, 437b chromium, 451 description of, 442–444, 448 prevention of, 465 zinc, 451 Galvanized steel, 449 Gamma decay, 479 Gamma rays, 476b–477b, 477, 496 Gas(es), 5b, 125–157. See also Gas laws atomic structure in, 6, 7 chemical reactions of, 139–141 gas phase equilibria, 399–400 ideal gas law, 129, 132, 136, 141, 152 kinetic–molecular theory, 141b, 141–148, 152, 155q, 366, 373 LeChatelier’s principle, 410–413 Maxwell-Boltzmann distribution, 143b, 143–144, 324, 325, 367 partial pressure, 136–139, 153q–154q phase changes, 296t, 296–298, 297 pressure and, 129–131, 129–132 properties, 128–129 sensors for, 148–151, 148–151, 156q state of, in writing chemical equations, 67 stoichiometry of reactions involving, 139–141, 152, 154q–155q STP conditions, 140–141 universal gas constant, 129b, 146 van der Waals constants, 146, 146t Gas furnaces, 313 Gas laws, 132–136, 151, 153q Avogadro’s law, 132 Boyle’s law, 132, 152 Charles’s law, 132, 152 Dalton’s law of partial pressures, 137 ideal, 129, 132, 136, 141, 152 Gas phase equilibria, 399–400 Gas sensors, 148–151, 148–151, 156q Gasoline, 100, 117. See also Fuel(s) additives to, 110–111, 117–118, 122q combustion of, 102, 107, 122, 287, 321 composition of, 100, 102, 103 energy density, 308t incomplete combustion of, 102, 103, 343 leaded, 117 octane, 102, 117 questions regarding, 119q reformulated, 117b seasonal formulations, 117 unleaded, 117 GDP. See Gross domestic product Geiger counter, 497b, 497 Germanium, 210, 254 Gibbs, J. Willard, 330 Gibbs free energy, 330b, 330–333, 339, 342q–343q Giga-, 14t Gold, 36t, 463 Goldschmidt process, 124 Graft copolymer, 270, 271b Grain elevator explosions, 385 Graphite, 242, 242 crystal structure of, 247, 249, 249 heat of formation, 303 intermolecular forces, 260–261, 261 phase change, 314 uses, 242

Gravity separators, 91 Green chemistry, 91b, 91–92, 96q, 113 Gross domestic product, 281, 281 Ground state, 172b Group 1 cations, 74t Group 1 metals, 48 Group (periodic table), 44, 44b, 46 Guldberg, Cato Maximilian, 398

H Haber process, 383 Half-life, 364–366, 482b, 502q Half-reaction equilibrium, 442 Half-reactions, 439b, 445–446, 467 Halides, 55t Halite, 85 Hall, Charles, 11, 459 Hall–Heroult process, 459, 459 Halogen lamps, 181, 181, 187 Halogens, 45b, 55, 187 Harmonic, 176 hcp. See Hexagonal close-packing HDPE. See High-density polyethylene Heat, 285b, 289–294. See also Thermodynamics calorimetry, 289b, 293, 293, 293–294, 313q chemical reactions and, 299–300 electricity production and, 298 enthalpy, 295–301, 314q indicating need for in writing chemical equations, 67 second law of thermodynamics, 326–327, 339, 341q specific, 289–290 work and, 285 Heat capacity, 290, 313q Heat flow, 312 enthalpy, 295–296 phase changes, 296t, 296–298, 297 Heat lamps, 165 Heat of formation, 300 Heat of fusion, 296, 296t Heat of reaction, 299b, 300, 314q Heat of vaporization, 296, 296t Hecto-, 14t Heisenberg, Werner, 180 Helium ionization energy, 190t van der Waals constant, 146t Helium balloons, 134, 151, 153 Helium-neon laser, 197q Helmholtz free energy, 330, 342 Hematite, 123, 450 Hepta-, 53t Heptane, 101t Heroult, Paul, 459 Hess’s law, 301–305, 314q, 333 Heterogeneous catalysts, 377, 377b Heterogeneous equilibrium, 400b Hexa-, 53t Hexagonal close-packing, 245, 245, 247 Hexane, 101t High-density polyethylene, 57, 57b, 58, 267, 337t High-level nuclear waste, 493 Homogeneous catalysts, 377b Homogeneous equilibrium, 400b Humidity, 126 Hund’s rule, 184b

Hybrid orbitals, 224–226, 225t, 225, 238q Hydrates, 41 Hydrazine molecular structure of, 89 production of, 115 synthesis of, 89 uses of, 89 Hydriodic acid, 78t Hydrobromic acid, 78t Hydrocarbons, 52b, 100b. See also Gasoline combustion, 68, 68, 69–70, 100, 102, 104, 209, 285 petroleum refining, 119 Hydrochloric acid, 78t aluminum hydroxide reaction with, 121 dilution calculation, 90 zinc metal reaction with, 153 Hydrocyanic acid, 78t, 422t Hydrofluoric acid, 78t, 422t, 433 Hydrogen combining with other elements, 45, 46, 213 combustion of, 108 covalent bond with, 50 energy density, 308t fluorine reaction with, 213, 213 as fuel, 106 ionization energy, 190t oxygen reaction with, 67, 67–68, 81, 121 spectrum of, 170, 170 van der Waals constant, 146t Hydrogen atom atomic spectra, 170 electron configuration, 184 orbital overlap, 222, 222 orbitals, 179t, 182, 183 Hydrogen azide, 239 Hydrogen bonding, 258–261, 275 Hydrogen carbonate, 55t Hydrogen chloride gas, 421 Hydrogen cyanide, 97 Andrussow process, 403, 409 boiling point, 263t industrial production of, 403, 409, 413 Hydrogen fluoride bond energy, 209 hydrogen bonding, 259 molecular structure of, 237 van der Waals constant, 146, 146t Hydrogen gas industrial production of, 121, 124 interstellar hydrogen clouds, 155, 155 iodine gas reaction with, 406–408 Hydrogen peroxide, 384, 386–388 Hydronium ion, 55t, 79, 420 Hydroxide anion, 55t Hydroxide ions, 79, 402 Hydroxides, 74t Hydroxyl radical, 496 Hypo-, 55 Hypochlorite ion, 55t Hypothesis, 11

I Ice melting of, 297, 331–332, 340 specific heat, 290t Ideal gas equation, 146–147

Ideal gas law, 129, 132, 136, 141, 152 Impurities, in metals, 9–10, 15 Incandescent lights, 159, 173, 195q, 197q–198q Indicators, 115, 115b Induced dipole, 257, 257 Inductive reasoning, 11b, 26 Inert gases, 45 Infrared radiation, 164, 164–165 Initiators, 41 Inorganic chemistry, 47b, 48–49, 60, 61q–62q Insoluble, 74b Instant ice pack, 340 Instantaneous dipole, 256–257 Instantaneous dipole–induced dipole forces, 256 Instantaneous rate of reaction, 352–353, 353 Insulators, 252b Integrated rate laws, 358b, 358–366, 382 for first-order reaction, 360–361 questions regarding, 385q–387q for second-order reaction, 362–364 for zero-order reaction, 359 Intermolecular forces, 256b, 256–261, 257–258, 275 dipole–dipole forces, 258, 258 dispersion forces, 256b, 256–258, 257, 259 hydrogen bonding, 258–261 questions regarding, 276q–277q Internal energy, 284b International Space Station, 466 International System of Units. See SI system Iodides, 74t Iodine gas, 406–408 Ion(s), 38b Coulomb’s law, 38b, 38–39 description of, 38–40, 59 in polymer chemistry, 40 properties, 39–40 questions regarding, 61q Ionic bonding, 42, 42, 235, 236q anion formation, 204–207, 205–206, 235 cation formation, 202–203, 235 Ionic compounds chemical nomenclature, 54–56, 55t formula unit, 42 Ionic lattice, 206, 206, 235 Ionic solids, 76 Ionization acid ionization constant, 422b, 422t base ionization constant, 422b periodic table and, 203 Ionization energy description of, 202 periodic table and, 188–190, 189, 190t Ionization gauge, 149–150, 150 Ionizing power, 495b, 495–497, 505q Ionizing radiation, 495–496, 505q Iridium, 276 Iron, 46, 276 atomic symbol, 36, 36t cations, 48, 49, 203 crystal structure, 247 electrochemical cell, 447–448 refining, 450 rust, 437, 443, 443, 451 Iron metal, 123 Iron ore, 97

Index

571

Iron(II) chloride, 55 Iron(II) ion, 55t Iron(III) chloride, 49, 49, 55 Iron(III) ion, 55t Irreversible reaction, 333b Isobutane, 263t Isobutene, 110 Isomers, 49, 100 Isotactic polymers, 268b Isothermal processes, 322 Isotopes, 34b, 34–35, 484t Isotopic abundance, 35b -ite (suffix), 55 ITER, 495

J Joule (unit), 285–286

K K capture, 480b “K shell,” 479 Kelvin temperature, 14t, 15, 15, 132, 135, 152 Kelvin (unit), 14t, 132, 135, 152 Ketones, 53t Kilo-, 13, 14t Kilogram (unit), 14t Kilowatt-hour (unit), 462 Kinetic energy, 142, 284b Kinetic–molecular theory of gases, 141b, 141–148, 145, 152, 155q, 366, 373 Kinetics. See Chemical kinetics Knocking (engine), 117 kPa (unit), 131 Kroto, Harry, 273 Krypton, 185 Krypton gas, 185

L Lambert, Frank, 340 Lanthanides, 46b, 186 Lasers, 167, 173, 192–193, 197q–198q Lattice energy, 207b, 236 Lattices, 42, 42b, 206, 206, 235–236 Law of conservation of matter, 68 Law of mass action, 398 Law of partial pressures, 137 Laws, 12b Laws of thermodynamics first law, 287, 312 second law, 326–327, 339, 341q third law, 327–329, 328t, 328, 342q LDPE. See Low-density polyethylene Lead atomic symbol, 36t fuel additives containing, 117 remediation of, 91 Lead azide, 83–84, 91 Lead iodide, 80 Lead nitrate potassium iodide reaction with, 80 sodium sulfate reaction with, 80 Lead-acid storage batteries, 456, 456b, 457 Leaded gas, 117 LeChatelier, Henri Louis, 410 LeChatelier’s principle, 410b, 410–415, 429, 432q LEDs. See Light-emitting diodes Lewis, G. N., 210

572

Index

Lewis dot symbol, 210b Lewis structures, 210–211, 211b, 215–221, 236 questions regarding, 237q–238q resonance hybrid, 221 VSEPR theory, 226, 231–233, 236 Light(s) artificial sources of. See Light sources coherent, 193b earth at night, satellite photo, 158 electromagnetic spectrum, 161, 161–169, 195q–196q frequency/wavelength relationship, 162 indicating need for in writing chemical equations, 67–68 monochromatic, 192b particulate nature, 165, 165–169, 166 photoelectric effect, 165b, 165, 165–167, 166, 198q photons, 167b, 167–169 properties, 163 refraction, 163, 163b, 163 speed of, 162 visible, 161b wave nature, 161–165, 162, 163 wave-particle duality, 167, 167b Light intensity, 14t Light sources fluorescent bulbs, 160, 160–161, 169–170, 173, 185, 195q, 197q–198q halogen lamps, 181, 181, 187 incandescent bulbs, 159, 160, 173, 195q, 197q–198q lasers, 167, 173, 192–193, 197q–198q light-emitting diodes, 173, 192–193, 193–194, 197q–198q organic light emitting diodes, 194, 274 types, 165 Light waves, 161, 162, 163 Light-emitting diodes, 173, 192–193, 193–194, 197q–198q Limestone, 97, 341 Limiting reactants, 108b, 108–112, 120q–121q Line structure, 49b, 50 Linear molecule, 227t, 230t Linear polyethylene, 57, 58 Liquid(s), 5b, 5–6, 261–265, 277q atomic structure in, 6, 7 boiling point, 263t, 263–264, 264 indicating state in writing chemical equations, 67 phase changes, 296t, 296–298, 297 surface tension, 264–265, 265 vapor pressure, 261b, 261–263, 262, 265 Liquid oxygen, 156 Liquid-liquid interactions, 265 Liquid–solid interactions, 265 Lithium band diagram, 251, 252 electron configuration, 184 ionization energy, 190t Lithium atoms, 250–251, 251 Lithium batteries, 310t Lithium hydroxide, 78t Lithium ion batteries, 310t London forces, 256. See also Dispersion forces Lone pairs, 211b Low-density polyethylene, 57, 57b, 58, 267, 337t

Lowry, Thomas, 420 Lubricants, 261

M Macromolecules, 32 Macroscopic perspective, 4–6, 81 Madelung constant, 206 Magic numbers, 488–489, 489b Magnesium, 96 combustion, 154 ionization, 203 ionization energy, 190, 190t Magnesium chloride, 342 Magnesium ion, 55t, 402 Magnetic quantum number, 177b, 178t Magnetic resonance imaging, 181 Magnetite, 450 Main group elements, 46b, 186, 210 Malleability, 5b, 249b, 250 Manganese, 276 Manometer, 148 Mass, 87 conversion to moles, 84–85 molar, 82, 82b, 84–90 units of, 14t Mass action, law of, 398 Mass action expression, 398 Mass defect, 487b Mass density, 5, 20–21 Mass number, 34, 34b Mass spectrometer, 38, 86, 151 Mass spectrum, 34–35 Matches, 110 Material selection, 24–25 Materials science. See also Polymer(s) buckminsterfullerene, 241, 273 fullerenes, 241, 273 invention of new materials, 272–274, 277q–278q nanotubes, 242, 242, 250–251, 251 superconductivity, 273, 273b Matter, 4b chemical properties, 5b conversion from raw material to waste, 3, 3 description of, 4–6 law of conservation of, 68 phases, 5b physical properties, 5b radiation and, interaction between, 495–498, 505q Matter–antimatter, 480 Maxwell-Boltzmann distribution, 143b, 143–144, 324, 325, 367 Mean free path, 145b Measurement, 27q–28q significant figures, 16–18 units, 13, 13–16, 14t, 15, 26 Mechanical energy, 285 Medical imaging, 498–499, 499, 505q Mega-, 14t Melting of ice, 297, 331–332, 340 Mendeleev, Dmitri, 44 Mercury, 60 atomic symbol, 36t in barometer, 130 electron configuration, 196, 198 in fluorescent lightbulb, 160, 169–170, 173

Mercury batteries, 454, 454b, 454 Mercury fulminate, 83–84 Mercury thermometer, 15 Mesoporous silica nanoparticles, 234, 234, 238q Metal(s), 43, 46b bonding in, 43 crystal structure, 247 electronegativity, 249 Fermi level, 252 impurities in, 9–10, 15 in periodic table, 46, 47, 59 physical properties, 249 Metal oxide semiconductor field effect transistor, 255 Metallic bonding, 43, 43b, 43, 249–251, 250–251, 276q band theory, 250b, 251–252, 275 sea of electrons model, 43, 250b, 250, 275–276 Metalloids, 46, 46b, 47 Meter (unit), 14t Methane, 101t in Andrussow process, 403, 409 combustion, 68, 104–105, 299–300, 302, 302 conversion to methanol, 426–427 energy density, 308t as fuel, 343 intermolecular forces, 260 Lewis structure of, 224 molecular structure of, 101t van der Waals constant, 146t Methyl cyanoacrylate, 97, 239 Methyl fluoride, 260 Methyl methacrylate, 49–50, 201, 211 Methyl tertiary-butyl ether. See MTBE Methylamine, 78t 2-Methylbutane, 100 Metric ton (unit), 24 Micro-, 14t Microscopic (particulate) perspective, 4b, 6–7, 81 Microstate, 324b Microwaves, 164, 164 Milli-, 14t Models, 11b, 11–12, 26, 27q Molar concentration, 89b, 93, 419 Molar heat capacity, 290, 290b Molar mass, 82, 82b, 84–90, 93, 105, 107 Molar solubility, 416b, 416–418 Molar volume, 140 Molarity, 88–89, 89b Mole(s), 14t, 82b, 82 conversion to mass, 85, 118 conversion to molecules, 84–85 description of, 82–84, 87, 93 energy of reaction, 305 questions regarding, 95q–96q Mole fraction, 137b Mole ratios, 83, 104b, 104–108, 118 Molecular equations, 77b, 93 Molecular formulas, 40b, 86 Molecular models, 226. See also Molecular shape VSEPR theory, 226, 229, 231–233, 236 Molecular scale engineering for drug delivery, 234–235, 238q Molecular shape, 226–233, 227t, 229t–230t, 233, 238q

Molecular structure, 59, 209–211 Molecularity, 374, 374t Molecules, 6b, 40b chemical formulas, 40b, 40–41, 67, 86 elemental analysis, 86b, 86–88 Monatomic ions, 38b, 38t, 39–40, 55 Mono-, 53t Monochromatic light, 192b Monomers, 32b Montreal Protocol, 377 MOSFET. See Metal oxide semiconductor field effect transistor Most probable speed, 144b MRI. See Magnetic resonance imaging MTBE, 110, 117–118, 118, 119 Multiplication, 17 Murphy’s law, 340

N Naming chemical compounds. See Chemical nomenclature Nano-, 14t Nanometer (unit), 20 Nanoparticles, 234, 238q Nanoscience, 1, 1 Nanotechnology, 241–242 Nanotubes, 242, 242b, 242, 250–251, 251 Naphthalene, 294 Natta, Giulio, 268 Natural gas, 343 Natural rubber, 272 n-Butane, 263t Neon, 190t Neon atom, 196 Neon lights, 164 Neopentane, 263t Nernst equation, 449b, 449–450, 467, 469 Net ionic equation, 77b, 81, 93, 94q–95q Neutralization, 79b Neutrino, 478, 480 Neutron, 33b, 34, 486 Neutron bombardment, 490, 495 Newlands, John, 44 Newton (unit), 130 Ni-cad batteries. See Nickel-cadmium batteries Nickel, 155 Nickel-cadmium batteries, 310t, 455, 455b, 455 Nickel-metal-hydride batteries, 310t, 456, 456b, 456 Nitrate anion, 55t Nitrates, 74t Nitric acid, 78t, 98 manufacture, 97, 120 reaction of copper with, 121 Nitric oxide, 306–307 Nitrites, 74t Nitrogen binary compounds with oxygen, 53, 54 electron configuration, 189 in explosives, 92 ionization energy, 190t van der Waals constant, 146t Nitrogen atom, 223 Nitrogen dioxide, 54 decomposition, 362–364 as pollutant, 127, 127–128

Nitrogen oxides, 54, 127 Nitroglycerin composition, 87 as explosive, 66, 66 molar mass, 83–84 molecular structure of, 66 Nitromethane, 85 Nobel, Alfred, 66, 66 Noble gases, 45, 45b Nodal plane, 179 Node, 179b Nomenclature, ions, 38 Nona-, 53t Nonane, 101t Nonattainment area, 127b Nonelectrolytes, 76 Nonmetals, 46b anion formation, 204 in periodic table, 46, 47, 59 valence electrons, 207 Normal boiling point, 263b Noryl, 51 Notation “ball and stick” model, 48, 229t–230t Lewis structures, 210–211, 215–221 line structure, 49b, 50 “space filling” model, 48 n-type semiconductors, 253, 253, 254b Nuclear charge, 182 Nuclear chemistry carbon dating, 475–476 cosmic rays, 475b, 475–476, 501q nuclear reactions, 475b, 476–477 radioactivity, 475b, 475–477, 477, 501q–502q Nuclear energy, 285 Nuclear fission, 488b, 504q Nuclear fusion, 488b, 494–495, 504q Nuclear mass, 487 Nuclear reactions, 68, 475b binding energy, 487–488, 488, 500 description of, 476–477, 487 equations for, 500 magic numbers, 488–489, 489b questions regarding, 501q–504q summary of, 500 Nuclear reactors, 492–493, 493, 501, 504q Nuclear shells, 488–489 Nuclear spin, 489 Nuclear stability, 485–487, 502q–503q Nuclear waste, 493 Nuclear Waste Policy Act, 493 Nucleus, 33b Nuclide, 476b Numbers questions regarding, 27q–28q rounding, 17 scientific notation, 16, 26 Nylon, 268, 269

O Observations description of, 26 interpreting, 10–11 questions regarding, 27q significant figures, 16–18 uncertainty in, 9–10

Index

573

Octa-, 53t Octahedral molecule, 227t, 230t Octane, 101t combustion, 102, 103, 120 energy density, 308t molecular structure of, 101t Octane number, 117 Octet rule, 210b, 211 !OH group, 52 OLED. See Organic light emitting diode Operator, 174b Optical disk drives, 196 Orbital(s), 173b, 178–181, 179–181, 182. See also Atomic structure antibonding molecular, 251b aufbau principle, 183b, 184, 186, 251 bonding molecular, 251b d, 179, 179 electron configurations, 182–183, 183 Hund’s rule, 184b hybrid, 224–226, 225t, 225, 238q p, 179, 179 potential energy and, 175 s, 179, 179 sigma bonds, 223, 223 visualizing, 178–181, 179–181 Orbital hybridization, 224–226, 225t, 225, 238q Orbital overlap, 238q chemical bonds, 221–224, 222b, 222–223 pi bond, 223 sigma bonds, 223, 223 Order of reaction, 354b, 354–355 Organic chemistry, 47b, 49–51, 59 notation, 48, 49b, 50 questions regarding, 61q–62q structural formula, 49 Organic compounds, 48, 52, 52b, 52t–53t Organic halides, 52t Organic light emitting diode, 194, 274 Osteoporosis, 505q Oxidation, 438, 467. See also Electrochemistry; Oxidation–reduction reactions Oxidation–reduction reactions, 308b, 438–440, 439, 467. See also Electrochemistry electrolysis, 7, 7, 458b, 458–464 Faraday constant, 449 Nernst equation, 449–450, 467, 469 Oxide ion, 204 Oxidizing agent, 439 Oxyanions, 55, 55t, 55b Oxygen in Andrussow process, 403, 409 ionization energy, 190t reaction of hydrogen with, 67, 67–68, 81, 121 van der Waals constant, 146t Oxygen difluoride, 215–216 Oxygenated fuel, 117 Ozone, 348 Chapman cycle, 376–377 decomposition, 350, 355, 357, 360, 365, 365–366, 369–371, 374 formation of, 127, 376 health hazards, 127 photodissociation in upper atmosphere, 360, 365, 365–366 resonance structure, 348

574

Index

in smog, 127 tropospheric, 379–381, 388q–389q Ozone alerts, 348 Ozone depletion, 348–350, 349, 360, 365, 365–366, 377, 382q Ozone hole, 350

P Packing efficiency, 243b, 243–245, 245–246, 275 Paint, 465 PAN. See Peroxyacetyl nitrate Partial pressure, 136–139, 153q–154q Dalton’s law of, 137 mass spectrometer, 151 Particles, wave-particle duality, 167, 167b Particulate matter, 127 Particulate (microscopic) perspective, 4b, 6–7, 81 Particulate nature of light, 165, 165–169, 166 parts per million (unit), 127b Pascal (unit), 130 Passivation, 465b Passive electrolysis, 458b, 458–460 Pauli exclusion principle, 182b, 197q Penetrating power, 495b, 495–497, 505q Penta-, 53t Pentane, 100, 100, 101t, 263t Peptide bonds, 47 Per-, 18, 55 Percentage yield, 113b, 121q Perchlorate ion, 55t Perchlorates, 74t Perchloric acid, 78t Period (periodic table), 44b, 44–46 Periodic law, 44b Periodic table, 36, 44–47, 59, 61q. See also Elements anion formation and, 204–205, 205 arrangement, 45–46, 47 atomic properties, trends in, 187–192, 197q atomic size and, 187, 187–188 crystal structure and, 247–248, 248 electron affinity and, 191, 191–192 electron configuration and, 185–187, 186, 197q electronegativity, 212, 212 ionization and, 203 ionization energy and, 188–190, 189, 190t views, 45, 47, 186, 212, 248 Periodicity, 44b Permanganate ion, 75 Permittivity, 39 Peroxyacetyl nitrate, 386 Pesticides, 124 PET. See Poly(ethylene terephthalate) Peta-, 14t Petroleum refining, 119 pH scale, 423, 423 Phase changes, 296t, 296–298, 297 heat of fusion, 296, 296t heat of vaporization, 296, 296t Phase diagram, 241, 241b, 242 Phases, 5b. See also Gases; Liquid(s); Phase changes; Solid(s) Phenols, 52t Phosgene, 120 Phosphate anion, 55t

Phosphate ions hydrogen ion reaction with, 404 Lewis structure of, 217–218 Phosphates, 74t Phosphoric acid, 78t, 122 Phosphorus, 254–255 forms of, 343 ionization energy, 190t Phosphorus pentachloride, 408 Phosphorus trichloride, 54, 408 Photochemical reaction, 68b, 127b Photochemical smog, 379 Photoelectric effect, 165b, 165, 165–167, 166, 198q Photon, 167b, 167–169, 479 Photon energy, 167 Photosynthesis, 342 Physical changes, 7 Physical properties, 5b, 26 Physical state of reactant, 67 Pi bonds, 223, 223b Picloram, 385 Pico-, 14t Picometer (unit), 204 Pigments, 272 Planck’s constant, 167 Plasmas, 5 Plasticizer, 272 Plastics recycling, 319, 319–320, 326–327, 336 advanced recycling, 327 bottles, 320 depolymerization, 327, 335 economics of, 335–338, 337t, 343q–344q entropy and, 326 questions regarding, 339q Plexiglas, 49, 321 PMMA. See Poly(methyl methacrylate) p-n junction, 255, 255b, 255 Polar bond, 213b, 235 Polar covalent bonding, 213, 213, 235 Polarity, 202b, 458, 459 Polarizability, 257b Pollutants, 97, 125, 127, 127–128. See also Air pollution gas pressure and, 132 removing, 140–141 Polonium-210, 276, 505q Polyatomic ions, 38b, 55, 55b Polycarbonates, 120 Polyethylene, 56–59 branched, 57 chemical formulas, 41 high-density, 57, 57b, 58, 267, 337t linear, 57, 58 low-density, 57b, 58, 58, 267, 337t melting point, 332–333 molecular structure of, 32, 32 production, 56–58 questions regarding, 62q–63q ultra-high molecular weight, 57–58 uses of, 32, 56 Poly(ethylene terephthalate), 337t industrial synthesis, 319, 319 recycling, 319–320, 335, 337t soft drink bottles, 319–320 Polyisobutylene, 123

Polymer(s), 31b, 31, 31–33, 48, 60q, 265–272, 277q addition, 266, 266–268, 267 additives to, 272, 428 atactic, 268b bonding in, 43 carbon-based, 47 chemical formulas, 41 condensation, 268b, 268–270 copolymers, 270b, 270, 270–271 cross-linking, 272 ions in formation of, 40 isotactic, 268b physical properties, 32, 271–272 silicon-based, 47 syndiotactic, 268b thermoplastic, 271, 271b, 271 thermosetting, 271, 271b, 271 Polymer backbone, 32, 32b Polymerization addition, 266, 266, 268 condensation, 268, 269 degree of, 267b entropy and, 325 thermodynamics, 329 Poly(methyl methacrylate) depolymerization, 327, 335 description of, 49–50 industrial synthesis, 321, 321, 327 methyl methacrylate polymerization to, 211 polymerization reaction, 201 Poly(phenylene oxide), 51 Polypropylene, 267, 337t Polystyrene, 50, 268, 271, 337t Polytetrafluoroethylene, 209, 209t Poly(vinyl alcohol), 218–219 Poly(vinyl chloride), 272, 337t chemical formula of, 41 molecular structure of, 32, 32–33 recycling of, 337t Poly(vinylidene chloride), 32, 32, 41 Poly(vinylpyrrolidone), 51 Portland cement, 392b, 398 Positron emission, 480–481 Positron emission tomography, 498–499, 499, 505q Potassium, 185 Potassium chlorate, 75 Potassium chromate, 76 Potassium hydroxide, 78t, 79–80 Potassium iodide, 80 Potassium ion, 55t Potassium permanganate decomposition, 124 solubility of, 75 Potential energy, 284b description of, 205, 208, 284 orbitals and, 175 Power plants, 298, 299 ppb (unit), 15b ppm (unit), 15b Precipitation, 80–81 Precipitation reaction, 80b Precision, 10, 10b Preexponential factor, 368b Prefixes for numbers, 53t in SI system, 14t

Pressure atmospheric pressure, 129, 129, 130 chemical equilibrium and changes in, 412–413, 413 gas laws, 132–136, 151, 153q gases and, 129–131, 129–132 measuring of, 130, 131 partial. See Partial pressure units, 130–132 Pressure sensors, 148–151, 148–151, 156q Pressure-volume work, 285b, 312 Primary cells, 309b, 309–310, 310t, 453b, 453–455 Primary standards of pollution, 127b Primers, for ammunition, 91 Primitive cubic, 245 Principal quantum number, 177, 177b, 178t Probability, 323, 323 Problem solving, 28q–29q conceptual problems, 21–22 ratios, 18–21, 26 Process engineering, 379 Product yield, 113–114 Products, 67b Propane, 101t boiling point of, 263t chemical bonds, 43 combustion of, 69–70, 119, 304, 387 energy density, 308t incomplete combustion of, 70–72 molecular structure of, 101t Propanoic acid, 422t Propellants, 377 Proteins, 47 Protons, 33, 33b, 34 PTFE. See Polytetrafluoroethylene p-type semiconductors, 254b, 254–255, 255 Purification, 327 Putrescine, 97 PV-work, 285b, 312 Pyridine, 97 Pyrrole, 97

Q Quadrillion, 282 Quality factor, 498b Quantum mechanical atomic model, 173–181, 196q, 256 Quantum numbers, 176b, 176, 176–178, 177t–178t, 182 Quantum theory aufbau principle, 183b, 184, 186, 251 Hund’s rule, 184b orbitals, 173b, 178–181, 179–181 Pauli exclusion principle, 182b, 197q quantum numbers, 176b, 176, 176–178, 177t–178t, 182 Schrödinger equation, 174b, 174–175, 177 terms defined, 174 uncertainty principle, 180b Quarks, 34, 486

R Radiant energy, 284 Radiation definition of, 476 detection methods, 497–498 dose measurements, 498, 498t

exposure, 498, 498t gamma, 479 imaging uses of, 498–499, 505q ionizing power of, 495b, 495–497, 505q matter and, interaction between, 495–498, 505q penetrating power of, 495b, 495–497, 505q Radio waves, 164 Radioactive decay description of, 365, 475b, 475–477 kinetics of, 481–484 questions regarding, 501q–502q Radioactive isotopes, 484t, 499–500 Radioactivity, 475b, 475–477, 477, 501q–502q Radiocarbon dating, 483–484 Radon, 496 Random copolymers, 270 Random error, 10b Rankine temperature scale, 16, 135 Rare gases, 45b Rate constant, 354b, 354–355, 380t, 383–384 Rate laws, 353–366, 382 chemical equilibrium and, 397–398 differential, 353–358, 354b, 382 integrated, 358–366, 382, 385q–387q questions regarding, 383q–385q Rate of condensation, 395 Rate of evaporation, 395 Rate-determining step, 376, 376b Rates of reaction, 350–356, 382, 382q–383q average rate, 352–353, 353 catalysis, 376–379, 388q differential rate law, 353–358, 354b, 382, 384q expressing of, 481 instantaneous rate, 352–353, 353 integrated rate laws, 358–366, 382, 385q–387q rate constant, 354b, 354–355, 383–384 rate laws, 353–366 rate-determining step, 376, 376b Ratios problem solving using, 18–21, 26 in stoichiometry, 104–108, 106, 118 RDX, 86–87 Reactants, 67b, 396. half-life, 364–366 limiting, 108b, 108–112, 120q–121q Reaction. See Chemical reactions Reaction mechanisms, 373b, 373–376, 387q–388q Reaction quotient, 398b Reaction rate, 350b, 351–352. See also Rates of reaction Reaction stoichiometry. See Stoichiometry Reactive intermediate, 374b Reactor core, 492b, 493 Recycling advanced, 327 depolymerization, 327, 335 economics of, 335–338, 337t, 343q–344q entropy and, 325–326 plastics, 319, 319–320, 327, 336, 339q Red phosphorus, 343 Redox reactions, 308b, 438b, 467. See also Oxidation–reduction reactions Reducing agent, 439

Index

575

Reduction, 438b, 467. See also Electrochemistry; Oxidation–reduction reactions Reformulated gasoline, 117b Refraction, 163, 163b, 163 Refrigerants, 124, 216, 288–289, 366, 377 Relative biological effectiveness, 498b Representative elements, 46b Resonance, 220b, 220–221, 221, 273–274 Resonance hybrid, 221 Resonance structures benzene, 221, 221 sulfur dioxide, 220–221 Reverse reactions, 394–396, 429 Reverse vending machines, 320 Reversible reaction, 333b RFG. See Reformulated gasoline Rocket engines, 111–112 Rocket fuel, 94, 111–112 decaborane, 118 diborane, 94 liquid oxygen, 156 Roller coaster, energy of, 284 Root-mean-square speed, 144b Rounding numbers, 17 Row (periodic table), 44–45 Rubber, 123, 272 Rust, 437, 443, 443, 451. See also Corrosion Rutherford, Ernest, 476, 504q

S Sacrificial anode, 466 Salicylic acid, 397 Salt. See Sodium chloride Salt bridge, 440b, 440, 440–441 Saturated solution, 75 sc lattice. See Simple cubic lattice Schrödinger equation, 174b, 174–175, 177 Scientific method, 4, 4b, 11–12 Scientific notation, 16, 16b Sea of electrons model, 43, 250b, 250, 275–276 Sea of instability, 485b, 485 Seaborg, Glenn, 503q Seaborgium, 503q Second ionization energy, 188 Second law of thermodynamics, 326–327, 339, 341q Second (unit), 14t, 461 Secondary batteries, 309b Secondary cells, 455b, 455–456 Secondary quantum number, 177, 177t, 177b, 178t Secondary standards of pollution, 127b Second-order integrated rate law, 362–364 Seesaw shape, 230t Semiconductors, 252–256, 253–255, 275–276 conductivity, 276 doping, 253, 253b, 253, 273–274 n-type, 253, 253 p-n junction, 255, 255 p-type, 254–255, 255 Semimetals, 46, 46b Shapes of molecules, 226–233, 227t, 229t–230t, 233 SHE. See Standard hydrogen electrode Shielding, 182b SI system, 13, 13–15, 14t

576

Index

Side reactions, 113b, 114 Sigma bonds, 223, 223b, 223 Significant figures, 16b, 16–18 Silane, 343 Silica, 234, 234 Silicon, 48 doping, 253, 253 heat of fusion, 314 ionization energy, 190t isotopes, 38 Lewis symbols, 210 as semiconductor, 252–253 Silicon compounds, 343 Silicon tetrachloride, 48, 48 Silicon-based polymers, 47 Silver, 36t, 55t Silver plating, 460, 460–461 Simple cubic lattice, 245, 247 Single bond, 211 Single event effect, 497 Slaked lime, 124 Smalley, Richard, 273 Smelting, aluminum, 23, 23–24 Smog, 125, 127, 379–380, 387 Soda cans. See Aluminum cans Sodium, 48 atomic symbol, 36t ionization energy, 190, 190t Sodium azide, 95 Sodium bicarbonate, 113 Sodium carbonate barium chloride reaction with, 80–81 Solvay process, 113 Sodium chloride crystal structure of, 43 dissolving of, 76 formula unit, 42, 43 ionic bonding in, 42, 42 solubility of, 75, 95 Sodium fluoride, 206–207 Sodium hydroxide, 78t, 116 Sodium hypochlorite, 89 Sodium ion, 38t, 39–40, 55t Sodium sulfate, 81 Solid(s), 5b atomic structure in, 6, 6, 6–7 bonding in, 249–256 description of, 275 indicating state in writing chemical equations, 67 packing efficiency, 245–246, 275 phase changes, 296t, 296–298, 297 questions regarding, 276q structure, 243–249, 245–246, 248 Solid fuel booster rockets, 111–112, 155 Solid structures coordination number, 247, 247b packing efficiency, 243–245, 245–246 Solid-state devices, 193 Solomon, Susan, 377 Solubility definition of, 74b, 416 rules for, 74t, 74–75 Solubility equilibria, 415–419, 432q–433q common ion effect, 418b, 418, 418–419 molar solubility, 416b, 416–418 Solubility product constant, 415–416, 416t

Soluble, 74b Solutes, 73b dilution, 90, 93 molarity, 88–89 Solutions, 72b concentration, 74b dilution, 90, 93 dilution formula, 90, 116 electrolytes, 75, 76 molarity, 88–89 saturation, 75 stoichiometry, 114, 114–116, 121q–122q Solvay process, 113 Solvents, 73b “Space filling” model, 48 Space program corrosion, 437, 466 decaborane as rocket fuel, 118 diborane as rocket fuel, 94 solid fuel booster rockets, 111–112 Sparingly soluble, 416b Specific heat, 289–290, 290b Specific heat capacity, 290b Spectator ions, 77b Spectra. See Atomic spectra; Electromagnetic spectrum Spectrometers, 34, 35 Speed, 142 average, 144, 144b Maxwell-Boltzmann distribution, 143b, 143–144, 324, 325, 367 most probable, 144b root-mean-square, 144b Speed of light, 162 Spherical polar coordinates, 179, 179 Spin, 489 Spin paired, 182b Spin quantum number, 181 “Spin up”/”spin down,” 182 Spitzer Space Telescope, 474 Spontaneous processes, 320–322, 321b, 339, 339q–340q Square planar molecule, 230t Square pyramidal molecule, 230t Standard Gibbs free energy change, 333b Standard hydrogen electrode, 444–445, 445, 468–469 Standard molar entropy, 328b, 328t Standard reduction potential, 445–449, 446t, 446b, 447, 468, 471 Standard state, 300, 441b Standard temperature and pressure, 140b, 140–141 State functions, 301, 301b, 322, 322 Statistical mechanics, 324b Steam, 290t Steel galvanized, 449 physical properties, 24, 24t rust, 437 Stiffness, 24 Stoichiometric coefficients, 69b, 404 Stoichiometry, 69b, 99b, 99–124 adjusting, 403–404 electrolysis and, 461–464, 472q energy and, 305–307, 314q equilibrium constant and, 399

fundamentals, 103–108, 119q–120q limiting reactants, 108b, 108–112, 120q–121q molar mass, 82, 82b, 84–90, 93, 95q–96q, 105 of reactions involving gases, 139–141, 152, 154q–155q percentage yields, 113–114, 121q rates of reaction and, 351–352 ratios, 26, 104–108, 106, 118 solutions, 114, 114–116, 121q–122q STP conditions, 140–141 theoretical yields, 113–114, 121q titration, 115b, 115–116, 116 word origin, 99 STP. See Standard temperature and pressure Stratosphere, 360, 364–365, 365, 377, 379 Strong acids, 78, 78t Strong bases, 78, 78t, 79 Strong electrolytes, 76 Strong force, 486, 486b Strontium hydroxide, 78t Structural formula, 49 Styrene, 50 Subshells, 177b Subtraction, 17 Sugar, 76, 76 Sulfate anion, 55t Sulfates, 74t, 75 Sulfides, 74t Sulfur electron configuration, 184 ionization energy, 190t Sulfur dioxide removal from industrial exhaust, 154 resonance structures, 220–221 van der Waals constant, 146t Sulfur hexafluoride, 121 Sulfur tetrafluoride, 231 Sulfur trioxide, 302–303 Sulfuric acid, 78t aluminum reaction with, 124 copper reaction with, 120 titration with sodium hydroxide, 116 Super glue, 97, 239 Superconductivity, 273, 273b Superheating, 263 Surface tension, 264–265, 265 Surroundings, 286b Symbolic representation, 8–9, 9 Syndiotactic polymers, 268b System, 286b Systematic error, 10, 10b, 10 Systeme International d’Units. See SI system

T Teflon, 209, 209t Temperature chemical equilibrium and, 414–415, 415t chemical kinetics and, 366–373, 367–369, 371, 372, 382, 387q gas laws, 132–136, 151 units, 14t, 15–16, 16, 135 Temperature scales, 15b, 132, 135 Tera-, 14t Termination, 266, 267 Termolecular steps, 374b Tetra-, 53t Tetraarsenic decaoxide, 120

Tetraethyl lead, 117, 118 Tetrahedral molecule, 227t, 229t Tetraphosphorus hexoxide, 54 Tetraphosphorus trisulfide, 107 Thallium-201, 506q Theoretical yield, 113b, 121q Theory, 11b, 26 Thermal energy, 285 Thermal pollution, 288 Thermochemical equation, 300, 300b, 305, 414 Thermocouple gauge, 149, 149–150 Thermodynamics, 191, 312, 339, 367. See also Heat delta, 287 endothermic, 296b, 299–300, 305, 322, 414–415 enthalpy, 295–301, 314q entropy, 322b, 322, 322–326, 323, 325, 328, 339, 340q–341q exothermic reactions, 296b, 299–300, 305, 322, 340, 414 first law, 287, 312 Gibbs free energy, 330–333, 339, 342q–343q heat of reaction, 299–300, 314q Hess’s law, 301–305, 314q, 333 of polymerization, 329 second law, 326–327, 339, 341q spontaneity, 320–322, 339, 339q–340q thermochemical equation, 300b, 305, 414 third law, 327–329, 328t, 328, 342q Thermolysis, 327 Thermometers, 15 Thermoplastic polymers, 271, 271b, 271 Thermosetting polymers, 271, 271b, 271 Third law of thermodynamics, 327–329, 328t, 328, 342q Three Mile Island nuclear accident, 310–311 Time, 14t Tin can, 443, 443 Tin coating, 464 Titanium, 24, 24t Titration, 115b, 115–116, 116 TNT. See Trinitrotoluene Torr (unit), 131b Torricelli, Evangelista, 131 Total ionic equation, 77b Transient dipole, 257, 257 Transition metals, 46b, 48–49, 49, 186, 203 Transmutation, 489, 489b, 504q Tri-, 53t Trichlorofluoromethane, 263t Trigonal bipyramidal molecule, 227t, 229t, 231 Trigonal planar molecule, 227t, 229t Trigonal pyramidal molecule, 229t Trinitrotoluene, 84 Triple bond, 211b Tritium, 494, 502q Tropospheric ozone, 379–381, 388q–389q T-shape molecule, 230t Tungsten electron configuration, 186 in halogen lamps, 181, 181 in incandescent lightbulb filament, 160, 186, 192 two-way arrow, in chemical equations, 78–79

U Ultra-high molecular weight polyethylene, 57–58 Ultraviolet radiation, 164, 164, 165, 239 Unbalanced chemical equations, 68 Uncertainty, in observations, 9–10 Uncertainty principle, 180b Uniform corrosion, 437b, 437 Unimolecular steps, 374b Unit cell, 245, 245b, 245 Units, 13, 13–16, 14t, 16, 26, 405 Universal gas constant, 129b, 146 Universe, 286b Unleaded gas, 117 Uranium-235, 484t, 490–492, 492 Uranium-238, 477, 484t, 503q Uranium enrichment, 97

V Vacuum, 130 Valence band, 252b Valence bond model, 222b Valence electrons, 185b atomic size and, 188 nonmetal elements, 207 Valence shell electron pair repulsion theory. See VSEPR theory van der Waals constant, 146t van der Waals equation, 146, 146b Vapor pressure, 261b, 261–263, 262, 265 Vaporization electricity production and, 298 heat of, 296, 296t Vehicle emissions, 97, 102 Vinylpyrrolidone, 51 Visible light, 161b Visualization, in chemistry, 22–24 Volatile organic chemicals, 127, 127b, 128, 379–380, 380t, 387 Volatility, 262, 262b Volcanoes, 137–138 Volta, Alessandro, 453 Voltmeter, 444, 444 VSEPR theory, 226, 226b, 231–233, 236 Vulcanization, 272

W Waage, Peter, 398 Waste energy, 288t, 288–289, 289 Water. See also Aqueous solutions boiling of, 7, 7 boiling point, 15, 16, 263t BrØnsted–Lowry theory, role in, 420–421 chemical bonds, 43 electrolysis, 7, 7 freezing point, 15, 16 mass density, 20 molar heat capacity, 290 molecular structure of, 7, 7 as solvent, 72–73 specific heat, 290t van der Waals constant, 146t vapor pressure, 262 in writing chemical equation, 67, 76 Watt (unit), 20, 462

Index

577

Wave(s) characteristics of, 163 diffraction of, 174b interference, 221–222, 222, 236 Schrödinger equation, 174b, 174–175, 177 Wave function, 174, 174b, 175–176, 178 Wave nature of light, 161–165, 162 Wavelength, 161b, 162, 167 Wave-particle duality, 167, 167b Weak acids, 78, 78t, 421–425 indicators, 115 reaction with strong base, 79 Weak bases, 78, 78t, 115, 421–425 Weak electrolytes, 76 Weathering, of concrete, 394, 394, 424 White phosphorus, 343

578

Index

Work, 285b, 312, 333 free energy and, 333 heat and, 285 second law of thermodynamics, 326–327, 339, 341q

X Xenon hexafluoride, 156 X-rays, 164, 164, 498–499

Y Yield, of chemical reactions, 113–114, 121q Yield strength, 24b, 24t Yocto-, 14t Yotta-, 14t Yucca Mountain storage facility, for nuclear waste, 493, 494, 504q–505q

Z Zepto-, 14t Zero, 16 Zero-order integrated rate law, 359, 359b Zetta-, 14t Ziegler, Karl, 268 Ziegler-Natta catalysts, 268 Zinc galvanic corrosion with chromium, 451 isotopes, 60 reaction with hydrochloric acid, 153 Zinc borate, 428 Zinc ion, 55t Zinc-air batteries, 310t, 454b, 454–455, 455, 471 Zinc-mercuric oxide cell, 454

SO ME USE FUL CONS TANT S (a more complete list appears in Appendix B) Atomic mass unit Avogadro’s number Electronic charge Faraday constant Gas constant

Pi Planck’s constant Speed of light (in vacuum)

1 amu 5 1.6606 3 10224 g N 5 6.02214179 3 1023 particles/mol e 5 1.60218 3 10219 coulombs F 5 96,485.3399 coulombs/mol e2 cal L atm R 5 0.08206 ———— 5 1.987 ———— mol K mol K J kPa dm3 5 8.314472 ———— 5 8.314472 ————— mol K mol K p 5 3.1415927 h 5 6.62606896 3 10234 J s c 5 2.99792458 3 108 m/s

SO ME USE F UL RE LATIONSHIPS Mass and Weight

Length

SI Base Unit: Kilogram (kg)

SI Base Unit: Meter (m)

1 kilogram 5 1000 grams 5 2.205 pounds 1 gram 5 1000 milligrams 1 pound 5 453.59 grams 1 amu 5 1.6606 3 10224 grams 1 gram 5 6.022 3 1023 amu 1 ton 5 2000 pounds

Volume

1 inch 5 2.54 centimeters (exactly) 1 meter 5 100 centimeters 5 39.37 inches 1 yard 5 0.9144 meter 1 mile 5 1.609 kilometers 1 kilometer 5 1000 meters 5 0.6215 mile 1 Ångstrom 5 1.0 3 10210 meters 5 1.0 3 1028 centimeters

Energy 3

SI Base Unit: Cubic Meter (m ) 1 liter 5 0.001 cubic meter 1 liter 5 1000 cubic centimeters 5 1000 mL 1 liter 5 1.056 quarts 1 quart 5 0.9463 liter 1 milliliter 5 0.001 liter 5 1 cubic centimeter cubic foot 5 7.475 gallons 5 28.316 liters 1 gallon 5 4 quarts

SI Base Unit: Joule (J) 1 calorie 5 4.184 joules 5 4.129 3 1022 L atm kg m2 1 joule 5 1 ——— 5 0.23901 calorie s2 1 joule 5 1 3 107 ergs 1 electron volt 5 1.6022 3 10219 joule 1 electron volt 5 96.485 kJ/mol 1 L atm 5 24.217 calories 5 101.325 joules

Pressure

Temperature

SI Base Unit: Pascal (Pa)

SI Base Unit: Kelvin (K)

kg 1 pascal 5 ———5 1 Newton/m2 m s2 1 atmosphere 5 760 torr 5 760 millimeters of mercury 5 1.01325 3 105 pascals 5 1.01325 bar 5 14.70 pounds per square inch 1 torr 5 1 millimeter of mercury

0 K 5 2273.15°C K 5 °C 1 273.15° °F 5 1.8(°C) 1 32° °F 2 32° °C 5 ————— 1.8°

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I NT ERNAT I O N A L TABLE OF ATOMIC WE IGHT S Atomic Number

Symbol

89 13 95 51 18 33 85 56 97 4 83 107 5 35 48 55 20 98 6 58 17 24 27 29 96 110 105 66 99 68 63 100 9 87 64 31 32 79 72 108 2 67 1 49 53 77 6 36 57 103 82 3 71 12 25 109 101 80 42

Ac Al Am Sb Ar As At Ba Bk Be Bi Bh B Br Cd Cs Ca Cf C Ce Cl Cr Co Cu Cm Ds Db Dy Es Er Eu Fm F Fr Gd Ga Ge Au Hf Hs He Ho H In I Ir Fe Kr La Lr Pb Li Lu Mg Mn Mt Md Hg Mo

Name Actinium Aluminum Americium Antimony Argon Arsenic Astatine Barium Berkelium Beryllium Bismuth Bohrium Boron Bromine Cadmium Cesium Calcium Californium Carbon Cerium Chlorine Chromium Cobalt Copper Curium Darmstadtium Dubnium Dysprosium Einsteinium Erbium Europium Fermium Fluorine Francium Gadolinium Gallium Germanium Gold Hafnium Hassium Helium Holmium Hydrogen Indium Iodine Iridium Iron Krypton Lanthanum Lawrencium Lead Lithium Lutetium Magnesium Manganese Meitnerium Mendelevium Mercury Molybdenum

Atomic Weight

Atomic Number

Symbol

[227] 26.9815386 [243] 121.760 39.948 74.92160 [210] 137.327 [247] 9.012182 208.98040 [272] 10.811 79.904 112.411 132.9054519 40.078 [251] 12.0107 140.116 35.453 51.9961 58.933195 63.546 [247] [281] [268] 162.500 [252] 167.259 151.964 [257] 18.9984032 [223] 157.25 69.723 72.64 196.966569 178.49 [270] 4.002602 164.93032 1.00794 114.818 126.90447 192.217 55.845 83.798 138.90547 [262] 207.2 6.941 174.9668 24.3050 54.938045 [276] [258] 200.59 95.96

60 10 3 28 41 7 102 76 8 46 15 78 94 84 19 59 61 91 88 86 75 45 111 37 44 104 62 21 106 34 14 47 11 38 16 73 43 52 65 81 90 69 50 22 74 112 116 118 115 114 113 92 23 54 70 39 30 40

Nd Ne Np Ni Nb N No Os O Pd P Pt Pu Po K Pr Pm Pa Ra Rn Re Rh Rg Rb Ru Rf Sm Sc Sg Se Si Ag Na Sr S Ta Tc Te Tb Tl Th Tm Sn Ti W Uub Uuh Uuo Uup Uuq Uut U V Xe Yb Y Zn Zr

Name Neodymium Neon Neptunium Nickel Niobium Nitrogen Nobelium Osmium Oxygen Palladium Phosphorus Platinum Plutonium Polonium Potassium Praseodymium Promethium Protactinium Radium Radon Rhenium Rhodium Roentgenium Rubidium Ruthenium Rutherfordium Samarium Scandium Seaborgium Selenium Silicon Silver Sodium Strontium Sulfur Tantalum Technetium Tellurium Terbium Thallium Thorium Thulium Tin Titanium Tungsten Ununbium Ununhexium Ununoctium Ununpentium Ununquadium Ununtrium Uranium Vanadium Xenon Ytterbium Yttrium Zinc Zirconium

Atomic Weight 144.242 20.1797 [237] 58.6934 92.90638 14.0067 [259] 190.23 15.9994 106.42 30.973762 195.084 [244] [209] 39.0983 140.90765 [145] 231.03588 [226] [222] 186.207 102.90550 [280] 85.4678 101.07 [267] 150.36 44.955912 [271] 78.96 28.0855 107.8682 22.98976928 87.62 32.065 180.94788 [98] 127.60 158.92535 204.3833 232.03806 168.93421 118.710 47.867 183.84 [285] [293] [294] [288] [289] [284] 238.02891 50.9415 131.293 173.054 88.90585 65.38 91.224

7

6

5

4

3

2

1

88

Ra (226)

87

Fr

(223)

(227)

Ac

89

138.9055

La

Y

39

44.9559

*

**

Note: Atomic masses are IUPC values (up to four decimal places). More accurate values for some elements are given on the facing page.

137.327

132.9055

57

56

87.62

85.4678

Ba

Sr

Rb

55

38

37

Cs

88.9059

40.078

39.0983

21

Sc

20

Ca

19

K

24.3050

22.9898

4

5

6

7

8

9

10 29

12

23

(268)

Db

105

180.9479

Ta

73

92.9064

Nb

41

50.9415

V

140.9076

140.116

24

U

92

144.24

Nd

60

(271)

Sg

106

183.84

W

74

95.96

Mo

42

51.9961

Cr

232.0381 231.0359 238.0289

91

Pa

90

Th

** Actinide Series

59

Pr

58

Ce

*Lanthanide Series

(267)

Rf

104

178.49

Hf

72

91.224

Zr

40

47.87

Ti

22

(237)

Np

93

(145)

Pm

61

(272)

Bh

107

186.207

Re

75

(98)

Tc

43

54.9380

Mn

25

(244)

Pu

94

150.36

Sm

62

(270)

Hs

108

190.2

Os

76

101.07

Ru

44

55.85

Fe

26

(243)

Am

95

151.964

Eu

63

(276)

Mt

109

192.22

Ir

77

102.9055

Rh

45

58.9332

Co

27

(247)

Cm

96

157.25

Gd

64

(281)

Ds

110

195.08

Pt

78

106.42

Pd

46

58.6934

Ni

28

(247)

Bk

97

158.9253

Tb

65

(280)

Rg

111

196.9665

Au

79

107.8682

Ag

47

63.546

Cu

(251)

Cf

98

162.50

Dy

66

(285)

Uub

112

200.59

Hg

80

112.411

Cd

48

65.38

Zn

30

13

12

Mg

11

9.0122

6.941

Na

10.811

Be

Li

(252)

Es

99

164.9303

Ho

67

(284)

Uut

113

204.3833

Tl

81

114.82

In

49

69.723

Ga

31

26.9815

Al

B

5

4

3

13

2

3

Metalloids

Nonmetals

1.0079

H

1

1

Metals

(257)

Fm

100

167.26

Er

68

(289)

Uug

114

207.2

Pb

82

118.710

Sn

50

72.64

Ge

32

28.0855

Si

14

12.011

C

6

14

(258)

Md

101

168.9342

Tm

69

(288)

Uup

115

208.9804

Bi

83

121.76

Sb

51

74.9216

As

33

30.9738

P

15

14.0067

N

7

15

(259)

No

102

173.054

Yb

70

(293)

Uuh

116

(209)

Po

84

127.60

Te

52

78.96

Se

34

32.065

S

16

15.9994

O

8

16

(262)

Lr

103

174.967

Lu

71

(210)

At

85

126.9045

I

53

79.904

Br

35

35.453

Cl

17

18.9984

F

9

1.0079

H

1

17

(294)

Uuo

118

(222)

Rn

86

131.29

Xe

54

83.80

Kr

36

39.948

Ar

18

20.1797

Ne

10

4.0026

He

2

18

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